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LSEh'A??Yof COWGFiESS 
Two Copies fie^ivwi 

JAJ^ 27 1905 

OOPY B. 



TABLE OF CONTENTS. 



CHAPTER I. 

INTRODUCTION. 



PAGE 

Natural Sciences — Physical and Chemical Changes — Chemical 
and Physical Properties — Energy — Matter — Mechanical 
Mixtures — Chemical Compounds — Elements 7 

CHAPTER n. 

OXYGEN. 

History — Occurrence — Preparation — Properties — Oxida- 
tion — Combustion — Kindling Temperature — Oxides — 
Ozone — Measurements of Gas Volumes 18 

CHAPTER HI. 

HYDROGEN. 

Historical — Occurrence — Preparation — Properties — Reduc- 
tion — Oxyhydrogen Blowpipe — Blast Lamp — Uses of 
Hydrogen SS 

CHAPTER IV. 

COMPOUNDS OF HYDROGEN AND OXYGEN — WATER AND HYDROGEN- 

DIOXIDE. 

Historical — Composition of Water — Occurrence — Impurities 
in Water — Purification of Water — Properties — Water of 
Crystallization — Uses of Water — Preparation, Properties 
and Uses of Hydrogen Dioxide. 46 

CHAPTER V. 

ATOMIC THEORY. 

Fundamental Laws of Matter — Atomic Theory — Value of a 

Theory — Atomic Weights 63 

CHAPTER VI. 

CHEMICAL EQUATIONS AND CALCULATIONS. 

Formulas — Equations — Kinds of Reactions — Chemical Equa- 
tions are Quantitative — Heat of Reaction 71 



4 Contents. 

CHAPTER VII. 

JiflTROGEN AND THE RARE ELEMENTS — ARGON — NEON — HELIUM — 
KRYPTON — XENON. 

PAGE 

Historical — Occurrence of Nitrogen — Preparation — Proper- 
ties — Rare Atmospheric Elements 79 

CHAPTER VIII. 

THE ATMOSPHERE. 

Constituents Essential to Life — The Function of Each — Con- 
stituents not Essential to Life — Quantitative Analysis of 
Air — Air a Mechanical Mixture — Changes in the Atmos- 
phere due to Respiration and Combustion — Properties of 
Ihe Atmosphere — Liquid Air 83 

CHAPTER IX. 

SOLUTIONS. 

Solution of Gases in Liquids — Solution of Solids in Liquids — 
Saturated Solutions — Crystallization — Properties of Solu- 
tions — Electrolytic Dissociation 91 

CHAPTER X. 

ACIDS, BASES, SALTS, VALENCE. 

Acids — Properties Common to Acids — Bases — Properties 
Common to Bases — Salts — Properties Common to Salts — 
Neutralization and the Dissociation Theory — Metals — Re- 
placing Power of the Metals — Valence — Structure of Com- 
pounds — Nomenclature of Acids, Bases and Salts — Acid, 
Basic and Normal Salts 100 

CHAPTER XI. 

COMPOUNDS OF NITROGEN. 

Decomposition of Organic Matter — Ammonia — Preparation of 
Ammonia — Properties of Ammonia — Volume Relation of 
Combining Gases — Nitric Acid — Preparation of Nitric 
Acid — Law of Mass Action — Properties of Nitric Acid — 
Nitrates — Nitrous Oxide — Preparation and Properties — 
Nitric Oxide — Preparation and Properties — Nitrogen Per- 
oxide — Preparation and Properties — Anhydrides 112 



Contents. 5 

CHAPTER XII. 

SULPHUR. 

PAGE 

Occurrence — Extraction — Physical Properties — Crystallog- 
raphy — Allotropic Forms of Sulphur — Chemical Proper- 
ties — Hydrosulphuric Acid — Preparation and Properties 

— Sulphur Dioxide — Preparation and Properties — Sulphur 
Trioxide — Preparation — Catalysis — Properties of Sulphur 
Trioxide — Sulphuric Acid — Manufacture — Properties — 
Sulphates — Thiosulphates — Monobasic and Dibasic Acids 

— Carbon Disulphide — Comparison of Sulphur and Oxy- 
gen — Selenium and Tellurium 126 

CHAPTER XIII. 

THE PERIODIC LAW. 

Natural Grouping of Elements — Metals and Non-Metals — 
Triad Families — Periodic Law — Periodic Table — Value 
of Periodic Law — Imperfections of the Law 144 

CHAPTER XIV. 

THE CHLORINE FAMILY. 

Fluorine — Occurrence — Preparation — Properties — Hydro- 
fluoric Acid — Preparation and Properties — Fluorides — 
Chlorine — Historical — Occurrence — Preparation — 
Properties — Bleaching — Nascent State — Hydrochloric 
Acid — Preparation and Properties — Aqua Regia — Oxy- 
gen Compounds ^ of Chlorine — Bromine — Occurrence — 
Preparation and Properties — Hydrobromic Acid — Prepa- 
ration and Properties — Oxygen Compounds — Iodine — Oc- 
currence — Preparation and Properties — Hydriodic Acid — 
Preparation and Properties 152 

CHAPTER XV. 

CARBON AND SOME OF ITS SIMPLER COMPOUNDS. 

Occurrence — Forms of Carbon — Properties of Carbon — Uses 

— Fuels — Hydrocarbons — Marsh Gas r— Acetylene — 
Other Hydrocarbons — Carbon Dioxide — Preparation — 
Properties — Carbonic Acid and the Carbonates — Carbon 
Dioxide and the Plant Life — Carbon Monoxide — Prepa- 
ration and Properties — Hydrocyanic Acid 169 



6 Contents. 

CHAPTER XVI. 

FLAMES — ILLUMINANTS. 

PAGE 

Flames — Supporter of Combustion — Appearance of Flames — 
Structure of Flames — Oxidizing and Reducing Flames — 
Luminosity of Flames — Bunsen Flame — Blowpipe Flame ' 

— Coal Gas — Water Gas — Natural Gas — Electric Fur- 
nace 183 

CHAPTER XVn. 

MOLECULAR WEIGHTS — ATOMIC WEIGHTS — FORMULAS. 

Determination of Equivalents — Relation of Equivalents to 
Atomic Weights — Combining Weights — Determination of 
Molecular Weights — Avogardro's Hypothesis — Hydrogen 
as Standard — Deduction of Atomic Weights — Raoult's 
Laws — Law of Dulong and Petit — Determination of 
Formulas 192 

CHAPTER XVHL 

THE NITROGEN FAMILY. 

Phosphorus — Occurrence — Preparation — Properties — Phos- 
phine — Oxides of Phosphorus — Acids of Phosphorus — 
Fertilizers — Arsenic — Occurrence — Preparation — Prop- 
erties — Arsine — Marsh's Test — Oxides of Arsenic — 
Acids of Arsenic — Sulphides of Arsenic — Antimony — 
Occurrence — Preparation — Properties — Stibine — Acids 
of Antimony — Sulphides of Antimony — Metallic Proper- 
ties of Antimony — Bismuth — Occurrence — Preparation 

— Properties — Compounds 203 

CHAPTER XIX. 

SILICON — BORON. 

Preparation of Silicon — Properties — Silicon Hydride — Sili- 
cides — Silicon Dioxide — Properties — Silicic Acids — 
Glass — Boron — Occurrence — Preparation and Proper- 
ties — Oxides — Acids — Borax 219 



INORGANIC CHEMISTRY, 



CHAPTER I. 
INTRODUCTION. 

The Natural Sciences. Before we advance very far 
in the study of nature, it becomes evident that the one large 
study must be divided into a number of more limited 
ones for the convenience of the investigator as well as of 
the student. These more limited studies are called the Nat- 
ural Sciences. 

Since these divisions are made for mere convenience, 
and not because there is any division in nature itself, it often 
happens that the different sciences are very intimately re- 
lated, and a thorough knowledge of any one of them in- 
volves a considerable acquaintance with several others. 
Thus the botanist must know something about animals as 
well as about plants ; the student of human physiology must 
know something about physics, as well as about zoology^ and 
anatomy. 

Mineralogy, Physics and Chemistry. Mineralogy, 
physics and chemistry form a group of sciences related in 
this close way, and it is not an easy matter to make an 
exact definition of any one of them. In a general way, 
However, it may be said that mineralogy deals with the 
composition and properties of the various substances which 
compose the earth's crust and with their distribution in 
nature. Chemistry and physics have for their object the 
study of the laws which govern the composition of all sub- 
stances, the properties of substances and the changes which 
they can be made to undergo. 
- (7) 



8 Inorganic Chemistry. ^ 

Two Kinds of Changes in Matter. An examination 
of these changes shows that they can all be classified under 
two heads : those in which no new kind of matter is formed, 
and those which result in the formation of new substances 
with new properties. The former are called physical 
changes, the latter chemical changes. 

1. Physical Changes. When a stick is broken, the 
pieces do not differ from the original stick save in size; 
the position of substances often changes, as when the leaves 
fall or trees are blown down ; a piece of iron may be mag- 
netized ; substances may be changed from cold to hot as the 
iron heated by the blacksmith in his forge, or the filament 
of the incandescent, lamp heated by the electrical current. 
In none of these changes is a new substance formed having 
a composition different from the substance undergoing the 

■ change. Such changes are called physical changes. 

2. Chemical Changes. In many other cases the sub- 
stances undergoing change lose their characteristic prop- 
erties and are converted into new substances. When a 
piece of coal or a stick of wood burns, it disappears as such, 
and in its place we have ashes and invisible gases. Gun- 

* powder, when ignited, explodes, and nothing but smoke and 
gases are found in its place. Sweet fruit juices ferment, 
and are changed into sour vinegar. Such changes are evi- 
dently quite different from the physical changes just men- 
tioned. New substances are formed in place of the ones 
undergoing change. These changes are called chemical 
changes. 

Changes in Physical State. It is not always easy to 
tell to which class a given change belongs, and many will 
require careful thought on the part of the student. One 
class of changes should be noted with especial care, since 
it is likely to prove misleading. It is a familiar fact that 
when water is cooled sufficiently it changes to ice ; when 
heated to the boiling point it changes to steam. Here we 



Introduction, 9 

have three different substances, water, ice and steam, one 
a liquid, one a soHd and one a gas, the properties of which 
differ widely. Apparently these are chemical changes. A 
careful study, however, shows that the composition of the 
three is just the same: they are all composed of the same 
substances, and in the same proportion. Many other solids 
can be melted into liquids without undergoing change in 
composition. Familiar instances are the melting of butter, 
wax and metals. The change of liquids into vapors is quite 
as common, but not so often noticed. In general almost any 
substance can exist in the three conditions, solid, liquid and 
gas, and we speak of them as the three physical states of 
matter. 

Chemical and Physical Properties. Many properties 
of a substance can be noted without causing the substance 
to undergo chemical change. Among these are its color, 
odor, taste, size, shape, weight, physical state. These are 
called its physical properties. Other properties are only dis- 
covered when the substance undergoes chemical change. 
These are called its chemical properties. Thus we find that 
coal burns in air, gunpowder explodes when fired, milk 
sours when exposed to air. 

Tv70 Factors in all Changes. In all the changes 
which matter can undergo, whether physical or chemical, 
two factors must be taken into account viz. energy and 
matter. 

I. Energy. It is a familiar fact that certain bodies 
have the power to do work. Thus water falling from a 
height upon a water wheel turns the wheel and in this 
way does the work of the mills. Magnetized iron at- 
tracts iron to itself and the motion of the iron as it 
moves towards the magnet can be turned into work. Coal 
on burning causes the engine to move and transports the 
loaded cars from place to place. When a body has this 
power to do work it is said to possess energy. 



lo Inorganic Chemistry. 

Law of Conservation of Energy. Careful experiments 
have shown that when one body parts with its energy, the 
energy is not destroyed but is transferred to another body 
or system of bodies. Just as energy can not be destroyed, 
neither can it be created. If a body gains a certain amount 
of energy, some other body has lost an equivalent amount. 
These facts are summed up in the law of conservation of 
energy which states that energy can neither he created nor 
destroyed, the amount of it in the universe being constant. 

Transformations of Energy. Although energy can 
neither be created nor destroyed, it is evident that it may 
assume many different forms. Thus the falling water may 
turn the electric generator and generate a current of 
electricity. The energy lost by the falling water is thus 
transformed into the energy of the electric current. This in 
turn may be changed into the energy of motion, as when 
the current is utilized for propelling the cars, or into the 
energy of heat and light, as when it is used for heating and 
lighting the cars. Again the energy of coal may be con- 
verted into energy of motion as when it is used as a fuel 
in steam engines. 

bince the energy possessed by coal only becomes avail- 
able when the coal is made to undergo a chemical change, 
it is sometimes called chemical energy. It is this form of 
energy with which we are especially interested in the study of 
chemistry. 

2. Matter. At first sight there appears to be no 
limit to the kinds of matter of which the world is made. 
But it has been found that all these varieties may be divided 
into three classes, viz., mechanical mixtures, chemical com- 
pounds and elements. These will now be briefly discussed. 

a. Mechanical Mixtures. If equal bulks of common 
salt and iron filings are thoroughly mixed together, a pro- 
duct is obtained, which, judging by its appearance, is a new 
substance. If it is examined more closelv however, it will 



Introduction. 



II 




Fig. 1 



be seen to be meiely a mixture of the salt and iron, each of 
v/hich substances retains its own peculiar properties. The 
mixture tastes just like salt; the iron particles can be seen, 
:^ and their gritty character detected. A magnet 
rubbed in the mixture draws out the iron just as 
if the salt were not there. The salt also can be sepa- 
rated from the iron quite easily. If 
several grams of the substance is 
placed in a test tube, and the tube 
half filled with water and thor- 
oughly shaken, the salt disappears, 
or dissolves in the water. The iron 
particles can then be filtered from 
the liquid by pouring the entire mass 
upon a piece of filter paper folded 
so as to fit into the interior of a fun- 
nel (Fig. i). The paper retains the 
solid but allows the clear liquid to drain through. The iron 
particles so obtained will be found to be identical with the 
original iron. The iron has undergone no change. The 
water of the salt solution can be removed by evap- 
oration. To accomplish this, the solution is poured 
into a small evaporating dish and gently heated 
(Fig. 2) until the water has all disappeared, or 
evaporated. The salt is left in the dish and on com- 
paring it with the original salt, it is found to be iden- 
tical with it. No new substance has 
been formed by mixing the salt and 
iron, and the mixture is called a 
mechanical mixture. 

Such mixtures are very common in 
nature. Thus in a piece of coarse 
grained granite, three quite different 




Fig. 2 



substances can be seen and picked out. Gold and sand 



12 Inorganic Chemistry. 

is a mechanical mixture, for the miner separates them by 
merely shaking the mixture in a pan, allowing running 
water to wash away the lighter sand, while the heavier 
gold sinks to the bottom of the pan. Evidently the com- 
position of such a mixture may vary greatly, and no two 
mixtures are likely to have the same composition. The 
ingredients of a mechanical mixture can often, though not 
always, be separated by mechanical means, such as sifting, 
sorting, magnetic attraction, or by dissolving one con- 
stituent and leaving the other undissolved. 

b. Chemical Compounds. If iron filings and pow- 
dered sulphur are thoroughly ground together in a mortar, 
'a yellowish green substance results. It might easily be 
taken to be a new compound ; but as in the case of the iron 
and salt, the ingredients can be readily separated. A magnet 
draws out the iron. Water does not dissolve the sulphur, 
but other liquids do, as for in- 
stance the liquid called carbon disul- 
phide. When treated with carbon 
disulphide, the iron is left' un- 
changed, and the sulphur can be ob- 
tained again, after filtering off the 
iron, by evaporating the liquid. The 
^■0 substance is therefore a mechanical 

^^ mixture. If now a new portion of 

^^s- 3 ^}^g mixture is placed in a dry test 

tube and carefully heated in the flame of the Bunsen burner, 
as shown in Fig. 3, a striking change takes place. The mix- 
ture begins to glow at some point, the glow rapidly ex- 
tending through the whole mass. If the test tube is broken 
and the product examined, it will be found to be a hard, 
black, brittle substance, in no way recalling the iron or the 
sulphur. The magnet no longer attracts it ; carbon disul- 
phid'e will not dissolve sulphur from it. It is a new sub- 




Introduction. 13 

stance with new properties, resulting from the chemical 
union of iron and sulphur. 

Again if a.^ small quantity of red oxide of mercury is 
heated in a hard glass test tube, it will be noticed that a 
bright silvery mirror gradually forms on the walls of the 
tube a little above the oxide. A glowing spark on a splint 
of wood, held at the mouth of the tube, bursts into flame, 
showing that a new, gaseous substance is being produced in 
the tube. On cooling the tube, the substance forming the 
silvery mirror does not change back again into the red 
oxide of mercury. It is a permanent new substance. By 
heating the red oxide, then, two new substances have been 
produced : the metalHc substance forming the silvery mir- 
ror, and the gas which caused the spark to burst into flame. 

\\'hen a substance contains two or more kinds of mat- 
ter, which are united chemically and cannot be separated by 
mechanical means, we say that it is a chemical compound. 
Thus the iron and the sulphur formed a chemical compound 
when they were heated together ; the red oxide of mercury 
is a chemical compound, because by chemical decomposi- 
tion two very different substances can be obtained from it. 
We shall see later that a chemical compound differs 
markedly from a mechanical mixture, in that its composi- 
tion is always exactly the same. 

Chemical Affinity the Cause of Chemical Combina- 
tion. The agency which causes substances to combine, 
and which holds them together when combined, is called 
chemical affinity. Yet the experiments just described show 
that heat is often necessary to bring about chemical action, 
and that it can also decompose some chemical compounds. 
The distinction between the cause producing chemical action, 
and the circumstances favoring it must be clearly made. 
Chemical affinity is always the cause of chemical union. 
Many agencies may make it possible for chemical affinity 



14 Inorganic Chemistry. 

to act by overcoming circumstances which stand in its way. 
Among these agencies are heat, electricity, and sometimes 
Hght. As a rule, solution also promotes action between two 
substances. Sometimes these agencies may overcome chem- 
ical attraction, and so occasion decomposition of a com- 
pound, as in the case of the decomposition of the red oxide 
of mercury by heat. 

c. Elements. It has been seen that strong heating of 
red oxide of mercury leads to the formation of two new 
substances. The question arises do these substances like- 
wise contain other substances ; that is are they also chem- 
ical compounds? All efforts have failed to get anything 
else from them but the metallic substance and the gas. 
Such substances are called elements. An element then is 
a substance from which no other substance can be obtained 
by any known means. It is not always easy to be certain 
that a given substance is really an element. Some way, 
as yet untried, may be successful in decomposing it into 
other simpler forms of matter, and the supposed element 
will then prove to be a compound. Water, lime and -many 
other familiar compounds were at one time thought to be 
elements. 

Alchemy. In olden times it was thought that some way 
could be found to change one element into another, and a 
great many efforts were made to accomplish this transforma- 
tion. Most of these efforts were directed toward changing 
the commoner metals into gold, and many fanciful ways for 
doing this were described. The chemists of that time were 
called Alchemists, and the art which they practiced was called 
Alchemy. The Alchemists gradually became convinced that 
the only way common metals could be changed into gold was 
by the wonderful power of a magic substance which they called 
the Philosopher's Stone, which would accomplish this trans- 
formation by its mere touch, and would in addition give per- 
petual youth to its fortunate possessor. No one has ever 



Introduction. 



15 



found such a stone, and no one has succeeded in changing 
one metal into another. 

Number of Elements. The number of substances 
now considered to be elements is not large — about eighty 
in all. Many of these are rare and very few of them make 
any large fraction of the materials in the earth's crust. 
Clark gives the following estimate of the composition of 
the earth's crust: 



Oxygen 

Silicon 27.-9% 

Aluminium 8.1% 

Iron 4-7% 



47.0% ;| Calcium 3.5% 

iMagnesium 2.5% 

Sodium 2-7% 

Potassium 2 . 4% 

Other elements.. 1.2% 



It is not necessary to study more than a third of the 
total number of elements to gain a very good knowledge of 
chemistry. 

Physical State of the Elements. About ten of the 
elements are gases at ordinary temperatures. Two — 
mercury and bromine — are liquids. The others are all 
solids, though their melting points vary through wide limits, 
from caesium which melts at 16° to elements which do not 
melt save in the intense heat of the electric furnace. 

Occurrence of the Elements. Comparatively few of 
the elements occur as uncombined substances in nature, 
most of them being found in the form of chemical com- 
pounds. When an elernent does occur by itself, such as 
gold, we say that it occurs in the free state or native ; when 
it is combined with other substances in the form of com- 
pounds, we say that it occurs in the combined state, or in 
combination. In the latter case there is usually little about 
the compound to suggest that the element is present in it; 
for we have seen that elements lose their owm peculiar prop- 
erties when thev enter into combination with other elements. 



1 6 Inorganic Chemistry. 

It would never be suspected that the red oxide of mercury 
contains a bright silvery metal, or that the eddish, earthy 
looking iron ore contains iron. 

Names of Elements. The names given to the ele- 
ments have been selected in a great many different ways. 
Some are very old and their original meaning is obscure. 
Such names are iron, gold, copper. Many indicate some 
striking physical property of the element. The name bro- 
mine is derived from a Greek word meaning a stench, re- 
ferring to the extremely unpleasant odor of the substance. 
The name iodine comes from a word meaning violet, allud- 
ing to the beautiful viokt color of iodine vapor. Some 
names indicate prominent chemical properties of the ele- 
ments. Nitrogen means the producer of nitre, nitrogen be- 
ing a constituent of nitre or saltpetre. Hydrogen means 
the producer of water, since the element forms water when 
burned. Others are named from mythology, as tantalum ; 
or from countries or localities, as germanium, and colum- 
bium. 

Symbols. In indicating the elements found. in com- 
pounds it is inconvenient to use such long names ; hence 
chemists have adopted a system of abbreviations. Some- 
times the initial letter of the name will suffice to indicate 
the element. Thus I stands for iodine, C for carbon. Usu- 
ally it is necessary to add some other characteristic letter of 
the name, since several names may begin with the same 
letter. Thus C stands for carbon, CI for chlorine, Cd for 
cadmium, Ce for cerium, Cb for columbium. Sometimes 
the symbol is an abbreviation of the old Latin name. In 
this way Fe (ferrum) indicates iron, Cu (cuprum) cop- 
per, Au (aurum) gold. These abbreviations are called sym- 
bols. They will become familiar through constant use. 

Law of Conservation of Matter. It has been seen 
tjiat matter often changes much in appearance ; elements 



Introduction. ly 

unite to form compounds which can be decomposed again 
into elements. In all such transformations it has been found 
that the amount of matter, or its mass, does not change at 
all. Xo matter is ever lost or created in these changes. 
This important truth is called the law of ':on5ervation of 
matter. The total amount of matter and of energy in the 
universe thus remains constant, though the form of each 
mav suffer many changes. 



CHAPTER 11. 
OXYGEN. 

History. The discovery of oxygen is generally at- 
tributed to the English chemist Priestley, who in 1774, ob- 
tained it by heating a compound of mercury "and oxygen 
known as mercuric oxide, the oxygen being expelled by the- 
heat. It is probable however that the Swedish chemist 
Scheele had previously obtained it although an account of 
his experiments Vv^as not published until 1775. The name 
oxygen signifies "to make acid." It was given to the ele- 
ment by the French chemist Lavoisier, since he believed 
that all acids owe their characteristic properties to its pres- 
ence — a view which we now know to be incorrect. 

Occurrence. Oxygen is by far the most abundant 
of all the elements. It occurs both in the free and in the 
combined state. 100 volumes of air contain about 21 vol- 
umes of free oxygen. In the combined states it forms eight- 
ninths of water and nearly one-half of the rocks com|)osing 
the earth's crust. It is also an important constituent of the 
compounds which compose plant and animal tissues. 

Preparation. Although oxygen occurs in the free 
state in the atmosphere, its separation from the nitrogen 
and other gases with which it is mixed is such a difficult 
matter that in the laboratory it has been found more con- 
venient to prepare it from its compounds. The most im- 
portant of the laboratory methods are the following: 

I. Preparation from Water. Water is a compound 
containing 11.18% of hydrogen and 88.82% of oxygen. 
It is easily separated into these constituents by passing a 
current of electricity through it under suitable conditions. 
The process will be described in the chapter on water. 

(18) 



Oxygen. 



19 



While this method of preparation is a simple one, it is not 
economical, on account of the cost of the electric current. 

2. Preparation from Mercuric Oxide. This method is 
of interest since it is the one which led to the discovery of 
oxygen. The oxide is placed in a small glass test tube and 
heated. The compound is thereby decomposed into mercury 
which collects on the sides of the glass tube forming a 
silvery .nirror, and oxygen which, being a gas, escapes from 
the tube. The presence of the oxygen is shown by lighting 
the end of a splint, extinguishing the flame and bringing 
the glowing coal into the mouth of the tube. The oxygen 
causes the glowing coal to burst into a flame. 

In a similar way, oxygen may be obtained from its com- 
pounds with some of the other elements. Thus manganese 
dioxide, a black compound of manganese and oxygen, when 
heated to a high temperature, loses one-third of its oxygen, 
while barium dioxide, on heating, loses one-half of its oxygen. 

3. Preparation from Potassium Chlorate. This is the 
method usually employed in the laboratory. Potassium, 
chlorate is a white solid composed of 31.9% potassium, 

28.9% chlorine and 39.1% oxygen. When heated it 
undergoes a series of changes in which the oxygen is 
finally all set free, leaving a compound of potassium 
and chlorine called potassium chloride. The change 
mav be represented as follows : 

potassium 

chlorine = 




oxygen 
(potassium 

chlorate) 



potassium 
chlorine 



+ oxygen 



(potassium 
chloride) 




Fig. 4 



The evolu- 
tion of the oxy- 
gen begins at 
about 400®. It 
has been found 
however that if 




20 Inorganic Chemistry. 

the potassium chlorate is mixed with about one-fourth its 
weight of manganese dioxide, the oxygen is given off at a 
much lower temperature. Just how the manganese dioxide 
brings about this result is not definitel)^ known. The amount 
of oxygen obtained from a given 
weight of potassium chlorate is ex- 
actly the same whether the manganese 
dioxide is present or not. So far as 
can be detected the manganese diox- 
ide undergoes no change. 
^^^- ^ The manner of preparing oxygen 

by this process is illustrated in the accompanying diagram 
(Fig. 4). A mixture composed of i part of manganese 
dioxide and 4 parts of potassium chlorate is placed in the 
flask A and gently heated. The oxygen is evolved and 
escapes through the tube B. It is collected by bringing over 
the end of the tube, the mouth of a bottle completely filled 
with water and inverted in a vessel of water as shown in the 
figure. The gas rises in the bottle and displaces the water. 
In the preparation of large quantities of oxygen, a 
copper retort, (Fig. 5) is often substituted for the glass flask. 

This method of collecting oxygen illustrates the general 
method used for collecting such gases as are insoluble in 
water or nearly so. The vessel C, (Fig. 4), containing the water 
in which the bottles are inverted, is called a "pneumatic trough." 

4. Commercial Methods of Preparation. Oxygen can now 
be purchased stored under great pressure in strong steel cyl- 
inders. This is prepared either by heating a mixture of potassium 
chlorate and manganese dioxide or by separating it from the 
nitrogen and other gases with which it is mixed in the atmosphere. 
Different methods have been employed for accomplishing this 
separation. The most important of these at the present time con- 
sists in liquifying the air by subjecting it to intense cold and great 
pressure. When the pressure is removed, the liquid nitrogen present 
evaporates first, leaving the liquid oxygen. As this evaporates the 
gas is pumped under high pressure into a strong steel cylinder 



Oxygen. 



21 




(Fig. 6) provided with a screw valve, so that the 
__^ oxygen may be drawn off at will. Oxygen so pre- 
pared is never entirely free from nitrogen, often 
containing as much as 15 to 20% of it. 

Physical Properties. Oxygen is a color- 
less, odorless, tasteless gas, slightly heavier than 
air. One liter of it weighs 1.428 grams, at a 
temperature of 0° and under a pressure of one 
atmosphere (1033 grams per square centimeter), 
while under similar conditions one liter of air 
weighs 1.292 grams. It is but slightly soluble in 
' water. As in the case of other gases, oxygen 
may be liquefied by applying very great pressure 
to the highly cooled gas. When the pressure is 
removed, the liquid oxygen rapidly boils, passing 
again into the gaseous state, since its boiling 



Fig. 



point under ordinary pressure is — 181.4°. 
Chemical Properties. At ordinary temperatures ox- 
ygen is not very active chemically. Alost substances are 
either not afifected by it at all, or the action is so slow as to 
escape notice. At higher temperatures however it is very 
active, forming compounds with all of the common elements 
except fluorine and bromine. This activity may be shown 
by heating various substances until just ignited and then 
bringing them into vessels of the gas, when they will burn 
with great brilliancy. Thus, as previously stated, a glow- 
ing splint if introduced into a jar of oxygen bursts into 
flame. Sulphur burns in the air giving a feeble light only ; 
in oxygen however the flame is increased in size and bright- 
ness. Substances like phosphorus which readily burn in 
air, burn in oxygen with dazzling brilliancy. Even sub- 
stances like iron which burn in air with great difficulty, 
readily burn in oxygen. 



22 Inorganic Chemistry. ^ 

The burning- of a substance in oxygen is due to the 
rapid combination of the substance or of the elements com- 
posing it with the oxygen. Thus when sulphur burns both 
the oxygen and sulphur disappear as such and there is 
formed a compound of the two which is an invisible gas, 
having the characteristic odor of burning sulphur. Simi- 
larly phosphorus on burning forms a white solid compound 
of phosphorus and oxygen, while iron forms a reddish 
black compound of iron and oxygen. 

Oxidation. The term oxidation is applied to the 
chemical change which takes place when a substance com- 
bines with oxygen. This process may take place rapidly as 
in the burning of phosphorus, or slowly as in the oxidation 
(or rusting) of iron when exposed to the air. It is always 
accompanied by the liberation of heat. The amount of heat 
liberated by tlie oxidation of a definite weight of any given 
substance is always the same, being entirely independent of 
the rapidity of the process. If the oxidation takes place 
slowly, the heat is generated slowly so that it is difficult to 
detect it. If the oxidation takes place rapidly however, the 
heat is generated in such a short interval of time that the 
substance may become white hot and even burst into a flame. 

Combustion. Kindling Temperature. When oxida- 
tion takes place so rapidly that the heat generated is suf- 
ficient to cause the substance to glow or burst into a flame 
the process is called combustion. In order that any sub- 
stance may undergo combustion it is necessary that it 
should be heated to a certain temperature known as the 
kindling temperature. This temperature varies widely for 
different bodies but is always definite for the same body. 
Thus the kindling temperature of phosphorus is far lower 
than that of iron but is definite for each. When any por- 
tion- of a substance is heated until it begins to burn, the com- 
bustion will continue without the further application of heat 



Oxygen. 23 

provided the heat generated by the process is sufficient to 
bring other parts of the substance to the kindHng temper- 
ature. On the other hand if the heat generated is not suf- 
ficient to maintain the kindhng temperature combustion 
ceases. 

Oxides. The compounds formed by the oxidation of 
any element belong to the class of compounds called oxides. 
Thus . in the combustion of sulphur, phosphorus and iron, 
the compounds formed are called respectively oxide of sul- 
phur, oxide of phosphorus and oxide of iron. In general 
then an oxide is a compound of oxygen with another ele- 
ment. A great many of these are known ; in fact the oxides 
of all the common elements have been prepared with the 
exception of those of fluorine and bromine. Some of these 
are familiar compounds. Water for exaniple is an oxide 
of hydrogen. 

Products of Combustion. The particular oxides 
formed by the combustion of any substance are called the 
products of combustion of that substance. Thus oxide of 
sulphur is the product of the combustion of sulphur ; oxide 
of iron is the product of the combustion of iron. It is evi- 
dent that the products of the combustion of any substance 
must weight more than the original substance, the increase 
in weight corresponding to the amount of oxygen combined. 
For example, when iron burns the oxide of iron formed 
weighs more than the iron which has undergone combus- 
tion. In some cases the products of combustion are invisible 
gases, so that the substance undergoing combustion is appa- 
rently destroyed. Thus when a candle burns, it is con- 
sumed and so far as the eye can judge, nothing is formed 
during combustion. That invisible gases are formed how- 
ever, and that the weight of these is greater than the weight 
of the candle may be shown by the following experiment. 



24 



Inorganic Chemistry. 



Two lamp chimneys are filled with the same amounts 
of the compoimd known as sodium hydroxide (caustic 
soda), and suspended from the beams of the balance as 
shown in Fig. 7. A piece of candle is placed on one ot 
the balance pans so that the wick comes just below 
the chimney, and th®- balance is brought to a level by adding 
weights to the other pan. The candle is then lighted. The 
products formed pass up through the chimney and are ab- 




Fig. 7 

sorbed by the sodium hydroxide. Although the candle 
burns away, the pan upon which it rests slowly sinks, show- 
ing that the combustion is attended by an increase in weight. 
Combustion in Air and in Oxygen. Combustion in 
air and in oxygen differ only in rapidity, the products 
formed being exactly the same. That the process should 
take place less rapidly in the former is readily understood, 
for .the air is only about one-fifth oxygen, the remaining 
four-fifths being inert gases. Not only is less oxygen avail- 



Oxygen. 2$ 

able therefore, but much of the heat is dissipated in raising" 
the temperature of the inert gases surrounding the sub- 
stance undergoing combustion. 

Combustion in the Broad Sense. According to the 
definition given above the presence of oxygen is necessary 
for combustion. The term is sometimes used however in a 
broader sense to designate any chemical change attended by 
the evolution of heat and light. Thus iron and sulphur, or 
hydrogen and chlorine under certain conditions will combine 
so rapidly that light is evolved. Whenever combustion takes 
place in the air however, the process is one of oxidation. 

Spontaneous Combustion. Some substances like 
phosphorus combine rapidly with oxygen even at ordinary 
temperatures. The heat generated by the process increases 
the rapidity of oxidation which in turn increases the amount 
of heat liberated so that the kindling temperature is soon 
reached and the substance bursts into a flame. This is an 
example of spontaneous combustion. Certain kinds of oils 
readily combine with oxygen at ordinary temperatures and 
not infrequently fires are traced to the presence of oily rags. 
Barns have been set on fire by the rapid oxidation of hay 
contained in them, while even heaps of coal have been 
known to undergo spontaneous combustion. 

The French chemist, Lavoisier (1743-1794), who gave to 
oxj'gen its name, was the first to show that combustion is 
due to union with oxj-gen. Previous to this, combustion was 
supposed to be due to the presence of a substance called phlo- 
giston. One substance was thought to be more combustible 
than another because it contained more phlogiston. When a 
substance burned, the phlogiston was supposed to escape. In 
accordance with this view, a substance on burning ought ta 
lose in weight. As has been shown, however, the opposite 
is true. 

Importance of Oxygen. Oxygen plays a very im- 
portant part in the processes of nature. Its presence in 



26 Inorganic Cheniistry. ^ 

the air is essential to all forms of life except certain minute 
forms of plant life. In respiration the air is taken into our 
lungs where a certain amount of the oxygen is absorbed by 
the blood. It is thus conveyed to all parts of the oody where 
it combines with the worn-out tissue, oxidizing it into com- 
pounds which ma}- readily be eliminated from the body. 
The heat generated by the oxidation is the source of the 
heat of the body. It was formerly thought that animals 
could not live in pure oxygen, on account of its great ac- 
tivity. Recent experiments seem to show however that this 
view is erroneous since the blood is capable of absorbing 
only a definite amount of oxygen ; hence any excess present 
in the lungs is without effect. 

The process of decay is also essentially one of oxida- 
tion. It will only take place however in the presence of 
certain minute forms of life known as bacteria. Just how 
these assist in the oxidation is not known. By this process 
the debris which collects on the surface of the earth- such 
as leaves and dead forms of animal life are slowly oxidized 
and changed into harmless products. Thus oxygen acts as a 
great purifying agent. 

Attention has been called to the fact that water dissolves 
a small amount of oxygen. All waters that come in contact 
with the air take up a certain amount of oxygen and this 
amount* although small is essential to all forms of aquatic 
life. Moreover the oxygen dissolved in water tends to 
purify it by oxidizing the impurities nearly always present 
in natural waters. 

Oxygen finds an important application in medicine in 
the treatment of such diseases as asthma where the patient 
is ujiable to inhale air sufficient to supply the necessary 
amount of oxygen. 



Oxygen. 2^ 



OZONE. 

When electric sparks are passed through oxygen it under- 
goes' a remarkable change, being converted into a new sub- 
stance called ozone. Ozone differs noticeably from oxygen in 
possessing a .characteristic odor, which is often perceived near 
electric machines. It also has different chemical properties, 
being much more active than oxygen. For example, oxygen 
has no action upon a sih^er coin, while ozone quickly colors it 
brown. 

Preparation. A convenient form of apparatus for the 
preparation of ozone in the laboratory is shown in Fig. 8. Two 
glass tubes are joined in the manner indicated and one surface of 

each coated with tin foil in 
•^ such a manner that an elec- 

jr trie discharge may be made 

3_^ to pass from one to the 
other by an electric machine 
or coil. A slow current of 



J^ 



Fig. 



oxygen is admitted at A, and passing between the two tubes, is 
subjected to the action of the electric discharge until it escapes at 
B. Only a small amount of the oxygen is changed into ozone 
by this treatment. The pure ozone may be obtained by cool- 
ing the mixture of oxygen and ozone escaping at B by means 
of liquid oxygen. Ozone being much more easily liquified than 
oxygen is thus changed to the liquid state, and by its evapora- 
tion the pure gas is obtained. Ozone is also formed in certain 
chemical processes. Thus if some pieces of phosphorus are 
placed in a bottle and partially covered with water, the presence 
of ozone may soon be detected in the air contained in the bottle. 
The Composition of Ozone. In the conversion of oxygen 
into ozone, it is possible to show by experiment that three 
volumes of oxygen form two volumes of ozone. The change 
moreover is a reversible one, the two volumes of ozone by 
heating being changed back into the three volumes of oxygen. 
In changing oxygen into ozone no other substance is added 
to it and of course none is removed, since it is an element. How 
then can we account for this change in properties? It can 
be shown that in the conversion of a given amount of oxygen 
into ozone a definite amount of heat is absorbed; moreover, 
if the ozone formed is converted again into oxygen, exactly 



28 . Inorganic Chemistry. ^ 

the same amount of heat is liberated. Oxygen differs from 
ozone therefore in its energy content, that of the ozone being 
greater, as indicated by the fact that more heat can be ob- 
tained from a given weight of it than from an equal weight of 
oxygen. The significance of this fact will be discussed in a 
later chapter. It is simply necessary here to note the fact 
that oxygen as well as a number of other elements exist in 
different forms or states, and that the conversion of the one 
state into another is always attended by the absorption or 
generation of heat. Such states are called allotropic states. 
Thus ozone is said to be an allotropic state of oxygen. 

Properties. Pure ozone is a gas having a bluish tint and 
characteristic odor. Its chemical properties are similar to those 
of oxygen, except that it is far more active. The pure gas 
is very explosive, breaking down into oxygen with the liber- 
ation of a large amount of heat. On account of its activity 
it is now prepared commercially and used in place of oxygen 
in a number of manufacturing processes. It is possible that 
traces of it may exist in the atmosphere, although its presence 
there has not been definitely established, the tests formerly 
used for its detection having been proved to be unreliable. 

Measurement of Gas Volumes. 

It is a well known fact that the volume of a gas is 
not constant, but varies with the temperature of the gas 
and the pressure to which it is subjected. It follows, 
therefore, that the weight of any definite volume of gas, say 
one liter, varies with the conditions of temperature and pres- 
sure. It has been agreed that the figures given to repre- 
sent the weight of a definite volume of any gas shall be its 
weight when the temperature of a gas is o° and when it is 
under a pressure equal to that exerted by the atmosphere at 
the sea level, viz: 1033 grams per square centimeter. These 
conditions of temperature and pressure are known as the 
standard conditions, and when the weight or volume of a gas 
is given, it is understood that the measurement was made under 
these conditions unless it is otherwise stated. Thus the weight 
of a liter of oxygen has been given as 1.428 grams, i. e., one 
liter' at a temperature of 0° and under a pressure of one atmos- 
phere, v/eighs 1.428 grams. 



Oxygen. 29 

Since the conditions of temperature and pressure under 
"which gases are measured in the laborator}- are different from 
the standard conditions, it becomes necessar}^ to be able to 
reduce the volume which a gas occupies under any given con- 
ditions of temperature and pressure to the volume it will occupy 
under standard conditions. ■ This may be done in accordance 
with the following laws. 

I. Variations in the Volume of a Gas Due to Change 
of Temperature. Law of Dalton and Gay Lussac. Increasing 
the temperature of a gas tends to increase its volume. The 
relation between the rise in temperature and in the increase 
'in volume was discovered independently by Dalton and by Gay 
Lussac in 1801. They were able to show by experiment the 
following fact or law: If a gas, having a temperature of o"" 
is heated, for a rise of each degree in temperature its volume 
is increased by 27-3 o± the volume which it occupied at o^ 
provided the pressure upon the gas remains unchanged. Simi- 
larly, if cooled below o^ its volume is decreased by 2-1-3 of 
the volume which it occupied at 0° for each degree through 
which the temperature has been lowered. If "\' represents the 
volume of gas at 0°, then the volume at 1° will be V+ irrV; 
at 2°, it will be V-f M^ ', or in general the volume, v, at the tem- 
perature t will be expressed by the formula 

<^* ^' = ''^^v 

or 

^''[ ^- = ^''1 + 2^3) 

Since arg- =^0.00366, the formula may be written 

(3) V = V (1 -f 0.00366t ) 

Since the value of V (volume under standard conditions) is 
the one generally sought, it is convenient to transpose the 
equation to the following form: 

(4) V= 

^ ' l-0.00366t 

The following problem will serve as an illustration: The 
volume of a gas at 20° is 750 cubic centimeters. Find the 
volume it will occupy at 0°, the pressure remaining constant. 



30 Inorganic Chemistry. 

In this case, v=750 c.c. and t^20. By substituting these 
values, equation (4) becomes 

^ = 1 + 0.00366 X 20 =^^^-^"" . 

2. Variations in the Volume of a Gas Due to Change of 
Pressure. Boyle's Law. The English scientist, Robert Boyle, 
as a result of his experiments, was the first to show that the 
volume of a gas varies inversely with the pressure to which 
it is subjected. This is known as Boyle's law. If the volume 
of a gas under a certain pressure is 4 liters, doubling the 
pressure will reduce the volume to 2 liters; or halving the 
pressure will increase it to 8 liters^ provided of course that 
the temperature remains constant. If V represent the volume 
when subjected to a pressure P and v represents its volume 
when the pressure is changed to p, then in accordance with 
the above law, V :v ::p :P or VP;=vp. In other words 
the product of the numbers representing the volume of a gas 
and the pressure to which it is subjected is a Constant. 

Since the pressure of the atmosphere at any point is in- 
dicated by the barometric reading, it is convenient to substi- 
tute the latter for the actual pressure, in the solution of prob- 
lems. The average reading of the barometer at the sea level 
is 760 millimeters, which corresponds to a pressure of 1033 
grams per square centimeter. The following problem will serve 
as an illustration: 

A gas occupies a volume of 500 c.c. in a laboratory where 
the barometric reading is 740 mm. What volume would it oc- 
cupy if the atmospheric pressure changed so that the reading 
became 750 mm.? 

Substituting the values in the equation VP=vp, we have 
500X740— vX7S0, or v=493.3 c.c. 

Inasmuch as corrections must as a rule be made for both 
temperature and pressure, it is convenient to combine the equa- 
tions given above for the corrections of each, so that the two 
corrections may be made in one operation. The following 
equation is thus obtained: 

v p 
^^^ ^"^760 (1 + 0.00366t) 
in which V represents the volume of a gas under standard con- 
ditions and V, the volume which it occupies when subjected to 
pressure p and temperature t. 



Oxygen. 31 

The following problem will serve to illustrate the appli- 
cation of this equation. 

A gas having a temperature of 20° occupies a volume of 
500 c.c. when subjected to a pressure indicated by a barometric 
reading of 740 mm. What volume would this gas occupy under 
standard conditions? 

In this problem v=:500, p=740 and t=:20. Substituting 
these in the above equation we get 

500X740 ^1-ofi 

760(1+0.00366X20) '^ -^ ^- ^• 

In many cases, gases are collected over water as ex- 
plained under the preparation of ox>:gen. In such cases there 
is present in the gas a certain amount of water vapor. This 
vapor exerts a definite pressure (called the tension of aqueous 
vapor) in opposition to the atmospheric pressure and hence 
must be subtracted from the latter in determining the effec- 
tive pressure upon the gas. The figures representing this 
tension at different temperatures are gi^^en at the end of this 
chapter. They express the tension in millimeters of mercury, 
corresponding to the barometric reading. 

In measuring the volume of a gas collected over water, 
the reading should be taken after raising or lowering the tube 
containing the gas until the level of the liquid inside and out- 
side" the tube is the same; for it is only under these conditions 
that the' gas is subjected to the atmospheric pressure as indi- 
cated b}^ the barometric reading. If the tube is not so adjusted, 
corrections must be made for the difference in pressure. The 
method of determining this correction may be illustrated by 
the following example: 

A tube graduated in cubic centimeters is partially filled 
with a gas collected over water, as shown in Fig. 9. The 
volume of the gas, as indicated by the reading on the graduated 
tube, is 250 c. c. Its temperature, as indicated by the reading 
of the thermometer placed near it, is 20°. The barometric reading 
is 740 mm. The level of ,the water inside the tube is found by 
measurement to be 100 mm. above the level outside. The prob- 
lem is to find the volume which this gas would occupy under 
standard conditions. 

It is evident that the pressure upon the gas is not indicated 
by the reading of the barometer, for the atmosphere pressing 



32 



Inorganic Chemistry. 



down on the surface of the liquid in the vessel B is not only 
„.-~^ compressing the gas in the tube A, but in 

addition is holding up a column of water 100 
mm. in height. The effective pressure upon the 
gas is equal therefore to the pressure as indi- 
cated by the barometric reading less the pres- 
sure spent in supporting the column of water. 
Since mercury is 13.6 times as heavy as water, 
a column of water 100 mm. in height would 
be equivalent to a column of mercury 7.3 mm. 
high. Subtracting this from the barometric 
reading we have 740 — 7-3^=722.7 mm. The 
gas being collected over water, the pressure 
is still further diminished by the tension of 
aqueous vapor. The value of this at 20° is 
- 17.4 m.m. Substracting this from 732.7, we get 
715.3 mm. as the actual pressure to which the 



Fig. 9 

gas is subjected. ,Now substituting in equation (5) given above, 
the values for v, t and p, we get 

250 X 715.3 



V=;^ 



760(1 -f 0.00366 X 20j 



= 219.2 c. c. 



Tension of Aqueous Vapor Expressed in Millimeters of Mercury 

Temp. 

16 



17 
18 

19 
20 



Pressure. 
.... 13 5 

14.4 

.... 15-4 
■.... 16.3 
.... 17.4 



Temp. 

21 ... 

22 ... 

23 ... 

24 ... 

25 ... 



Pressure. 
.... 18.5 
.... 19-7 
.... 20 . 9 
.... 22 . 2 
23.6 



CHAPTER III. 

HYDROGEN. 

Historical. Hydrogen was first clearly recognized as 
an element by the English investigator Cavendish, who, in 
1776, obtained it in a pure state, and showed it to be dif- 
ferent from the other inflammable airs or gases which had 
long been known. Lavoisier gave it the name hydrogen, 
signifying "water producer" since it had been found to be 
a constituent of water. 

Occurrence. In a free state hydrogen is found in the 
atmosphere but only in traces. Large quantities of it oc- 
cur however in the gaseous matter surrounding the sun 
and certain other stars. In the combined state it is widely 
distributed, being a constituent of water as already indi- 
•cated. Combined with carbon it occurs in crude petro- 
leum and natural gas. Nearly all of the products of the 
animal and vegetable kingdoms, such as sugar and starch, 
also contain it. 

Preparation from Water. 

a. By flic Electric Current. As already indicated in 
the preparation of oxygen, w^ater is easily separated into its 
constituents, hydrogen and oxygen by passing an electric 
current through it under certain conditions. 

b. By the Action of Metals. When brought in con- 
tact with certain metals under appropriate conditions, water 
gives up a portion or the whole of its hydrogen, its place 
being taken by a definite amount of the metal. In the case 
of the metal sodium, this change occurs at ordinary temper- 
atures. Thus if a bit of sodium is thrown on water an 

(33) 



34 



Inorganic Chemistry. 



action is seen to take place at once, sufficient heat being; 
generated to melt the sodium which runs about on the sur- 
face of the water. The change consists in the displace- 
ment of one-half of the hydrogen of the water by the so- 
dium, and may be represented as follows : 

r hydrogen f sodium 

sodium + I hydrogen = hydrogen + -<^ hydrogen 
(oxygen [oxygen 

(water) (sodium hydroxide) 

The sodium hydroxide formed is a white solid which 
remains dissolved in the undecomposed water and may be 
obtained by evaporating the solution to dryness. 




Fig. 10 



Other metals such as magnesium and iron decompose 
water rapidly, but only at higher temperatures. The 
method of effecting this decomposition and collecting the 
liberated hydrogen is illustrated in Fig. lo. A porcelain or 
iron tube, B, about 50 cm. in length and 2 or 3 cm. in 



Hydrogen. 35 

width is partially filled with fine iron wire or tacks, and 
connected as shown in the figure. 

The flask A is partially filled with water. The tube is 
heated, slowdy at first, until the iron is red hot. Steam is 
then conducted through the tube by boiling the water in the 
flask A. The hot iron combines with the oxygen in the 
steam setting free the hydrogen which is collected over 
water. The gas which passes over first is mixed with the 
air previously contained in the flask and tube, and is al- 
lowed to escape since a mixture of hydrogen and oxygen 
or air explodes with great violence zvhen brought in con- 
tact zvith a Ha me. It is evident that the flask A must be 
disconnected from the tube before the heat is withdrawn. 

That the gas obtained is different from air and oxygen 
may be shown by holding a bottle of it mouth downward 
and bringing a lighted splint into it. The hydrogen is ig- 
nited and burns with an almost colorless flame. 

If, at the close of the experiment the iron in the tube 
is examined, it will be found that at least a portion of it has 
been changed to a reddish black substance. This has been 
proved to be an oxide of iron, identical with that obtained 
by the combustion, of iron in oxygen. The change which 
has taken place may be represented as follows : 

-, ^ J ^ hydrog-en f iron 

iron+^ hydrogen = -, -^ -, ^ + < 

\ oxygen hydrogen ^ (oxygen 

( water ) ( iron oxide ) 

The iron has thus taken the place of the whole of the 
hydrogen in the water ; in other words it has combined 
with the oxygen and liberated the hydrogen. 

If the element carbon, in the form of charcoal or coke 
is substituted for the iron in the above experiment, the hot 



36 Inorganic Chemistry. 

carbon, like the iron combines with the oxygen and liberates the 
hydrogen. The change may be represented as follows: 

carbon + I hydrofen = ^Y^^ogen f carbon 
i oxygen hydrogen ^ I oxygen 

(water) (carbon monoxide) 

The carbon monoxide formed is a gas however, so that the 
product obtained by this method is not pure hydrogen but a 
mixture of two gases. Since both of these gases burn with a hot 
flame, the mixture is often manufactured on a large scale and used 
as a fuel under the name "water gas." 

While hydrogen may thus be obtained readily from 
water the methods are not economical and are therefore 
little used. 

Preparation from Acids. 

a. Usual Laboratory Method. Among the most 
important compounds with which the chemist has to deal 
are those included under the name, ''acids." Hydrogen is 
an essential constituent of all acids. Moreover when cer- 
tain metals are brought in contact with an acid, the hydro- 
gen contained in the acid is liberated, being replaced by a 
definite amount of the metal. Although this reaction is a 
general one, it has been found most convenient in prepar- 
ing hydrogen by this method to use for the acid, either 
hydrochloric or dilute sulphuric and for the metal either 
zinc or iron, generally the former. 

Both sulphuric acid and zinc, if impure, are likely to contain 
small amounts of arsenic. Such materials should not be used in 
preparing hydrogen since the arsenic present combines with a por- 
tion of the hydrogen to form a very poisonous gas known as arsine. 
On the other hand chemically pure sulphuric acid, i. e. sulphuric 
acid that is entirely free from impurities, will not act upon chem- 
ically pure zinc. The reaction may be started however by the 
addition of a few drops of a solution of copper sulphate. 

To prepare hydrogen by this method the metal is placed 
in a flask or wide mouth bottle A (Fig. 11) and the acid 



Hydrocreii. 



37 



slowly added through the funnel tube B. The metal dis^ 
solves in the acid. The hydrogen which is liberated es- 




capes through the exit tube C and is collected over water. 
It is evident that the hydrogen which passes over first is 
mixed with the air contained in the bottle ; hence care 
must be taken not to bring a flame near the exit tube, since 
as previously stated such a mixture explodes with great 
violence when brought in contact with a flame. 

The changes which take place when zinc and dilute sul- 
phuric acid are used may be represented as follows : 

C h3'drogen C zinc 

zinc -f -' sulphur =h3'drogen + - sulphur 
( 0X3' gen ( ox3'gen 

( sulphuric acid) (zinc sulphate) 

In other words the zinc has taken the place of the hy- 
drogen in sulphuric acid, which is a compound of hydrogen, 
sulphur and ox3-gen. The resulting compound contains zinc, 
sulphur and ox3'gen and is known as zinc sulphate. This 
is a white solid which is soluble in water and hence dissolves 
as fast as formed in the water present in the acid. It ma3^ 



38 Inorganic Chemistry. ^ 

be obtained by evaporating the liquid which is left in the 
flask A, Fig-. 11, after the metal has passed into solution. 

When zinc and hydrochloric acid are used the follow- 
ing changes take place. 

zinc -\- < 7, .^^ = hydrogen + \ ^Y\^ . 

(^chlorine -^ ^ ( chlorine 

(hydrochloric acid) (zinc chloride) 

The zinc chloride is a white solid, which dissolves in 
the water of the flask and may be obtained by evaporating 
the solution to dryness. When iron is used the changes 
which take place are exactly similar to those just given for 
zinc. 

Physical Properties. Hydrogen is similar to oxy- 
gen in that it is a colorless, tasteless, odorless gas. It 
is characterized by its extreme lightness, being the light- 
est of all known substances. One liter of it weighs 
only 0.08984 grams. On comparing this weight with 
that of an equal volume of oxygen viz, 1.428 grams, 
the latter is found to be 15.8 or 
in round numbers 16 times as heavy. 
Similarly air is found to be 14.4 times 
as heavy as hydrogen. Its extreme 
lightness may be shown by blowing 
soap bubbles with it, the bubbles rising 
in the air with great rapidity. This 
suggests its use for the inflation of bal- 
loons. On account of its lightness it is possible to pour it 
''upward" from one bottle into another. Thus if the bottle 
A, (Fig. 12) is filled with hydrogen, placed mouth down- 
ward by the side of bottle B filled with air, and then gradu- 
ally changed as indicated in the figure to an upright position 
the hydrogen will flow upward into bottle B displacing the 
air. ^ Its presence in bottle B may be then shown by bringing 
a lighted splint to the mouth of the bottle, the hydrogen be- 




Hydrogen. 39 

ing ignited by the flame. It is evident from this experiment 
that in order to retain the gas in an open bottle, the bottle 
must be held mouth downward. 

Hydrogen is far more difficult to liquify than any 
■other gas, with the exception of the rare element helium 
which has as yet resisted all efforts to liquify it. The Eng- 
lish scientist Dewar however in 1898 succeeded not only in 
obtaining it in liquid state but also in a solid state. Liquid 
hydrogen is colorless and has a density of only 0.07. Its 
boiling point under atmospheric pressure is — 252°. Under 
diminished pressure the boiling point is lowered to — 258°, 
a point only 15° above absolute zero. 

The solubility of hydrogen in water is very slight being 
still less than that of oxygen. 

Chemical Properties. At the ordinary temperature, 
hydrogen is an inactive element. At higher temperatures 
however it combines with a number of elements being re- 
markable for its affinity for oxygen. On account of this affin- 
ity, a mixture of hydrogen and oxygen or hydrogen and air, 
(which is about one-fifth oxygen) explodes with great vio- 
lence when heated to the kindling temperature. Neverthe- 
less under proper conditions, hydrogen may be made to burn 
quietly in either of these gases. To bring this about the 
hydrogen is generated in the bottle A, (Fig. 13) is dried by 
conducting it through the tube X, filled with some substance 
(generally calcium chloride) which has a great attraction for 
moisture, and escapes through the tube T, the end of which 
is drawn out to a jet. The hydrogen first liberated mixes 
with the air contained in the generator. If a flame is 
brought near the jet before this mixture has all escaped, a 
violent and very dangerous explosion results, since the en- 
tire apparatus is filled with the explosive mixture. On the 
other hand if the flame is not applied until all the air has 



40 



Inorganic Chemistry. 



been expelled the hydrogen is ignited and burns quietly, 
since only the small amount of it which escapes from the jet 
can come in contact with the oxygen of the air at any one 
time. 




Fig. 13 



Before igniting the hydrogen therefore it is absolutely neces- 
sary to know that it is free from air. This is determined by test- 
ing small amounts of the escaping gas. A convenient and safe 
method of doing this is to fill a test tube with the gas by bringing 
it down over the jet. The hydrogen on account of its lightness 
collects in the tube. After leaving it over the jet for a few mo- 
ments in order that it may be filled with the gas, the tube is 
gently carried mouth downward to the flame of a burner placed 
not nearer than an arm's length from the jet. If the hydrogen is 
mixed with air a slight explosion occurs ; but if pure, it burns 
quietly in the tube. The operation is repeated until the gas burns 
quietly, when the tube is quickly brought back over the jet for 
an instant, whereby the escaping hydrogefi is ignited by the 
flame in the tube. 

The hydrogen flame is almost colorless and very hot. 
The combustion of the gas is due, of course to its union 
with oxygen. The product of the combustion is therefore 



Hydrogen. 



41 



a compound of hydrogen and oxygen. That the compound 
formed is water which as already indicated is composed of 
these two elements, may be shown by holding over the flame 
a cold and dry bell jar or bottle (Fig. 13). The steam is 
condensed, the water collecting in drops on the sides of the 
jar. 

That a mixture of hydrogen .and air is explosive may 
be shown safely as follows. A cork through which passes 
a short glass tube about i cm. in 
^diameter is fitted air tight into the 
tubule of a bell jar of two or three 
liters capacity. (A thick glass bot- 
tle with bottom removed ma}^ be 
used.) The tube is closed with a 
small rubber stopper and the bell jar 
filled with hydrogen, the gas being 
collected over w^ater. When entirely 
filled with the gas the jar is removed 
from the water and supported by 
blocks of wood in order to leave the 
bottom of the jar open as shown in 
Fig. 14. The stopper is now removed 
from the tube in the cork and the hydrogen, which on account 
of its lightness escapes from the tube, is at once lighted. As the 
hydrogen escapes, the air flows in at the bottom of the jar 
and mixes with the remaining portion of the hydrogen, so 
that a mixture of the two soon forms and a loud explosion 
results. The explosion is not dangerous, since the bottom of 
the jar is open, thus leaving room for the expansion of the 
hot gas. 

Since air is only one-fifth oxygen the remainder being in- 
ert gases, it may readily be inferred that a mixture of hydrogen 
with pure oxygen would be far more explosive than a mixture 
of hydrogen with air. Such mixtures should not be made 
except in small quantities and by experienced teachers. 

Hydrogen does not Support Combustion. While hy- 
drogen is readily combustible it is not a supporter of com- 
bustion i. e., substances will not burn in it. This mav be 




Fig. 14 



42 



Inorganic Chemistry. 



mm 



shown by bringing a lighted candle supported 
by a stiff wire into a bottle or cylinder of the 
pure gas as shown in Fig. 15. The hydrogen is 
ignited by the flame of the candle and burns at 
the mouth of the bottle. When the candle is 
thrust up into the gas, its flame is extinguished on 
account of the absence of oxygen. If slowly with- 
drawn the candle is relighted as it passes through 
the layer of burning hydrogen. 

Reduction. On account of its great afflnity 
for oxygen, hydrogen has the power of abstract- 
ing it from many of its compounds. Thus if a 
stream of hydrogen, dried by passing through the 
tube B (Fig. 16), filled with calcium chloride, is conducted 
through the tube C containing some copper oxide, heated 
to a medium temperature, the following change takes place. 

hydroo:en+ \ ^^PP — copper + i •^ ^ 
^ ^ [ oxygen ^^ [ oxygen 

(copper oxide) (water) 



Fig. 15 







k 



B 



p=:= ^f?&M:Tiizmx>i d 1 t 



( ■ , ^ , . -J 




Fig. 16 



Hydrogen. 43 

The vrater formed collects in the cold portions of the 
tube C near its end. 

Jn this experiment the copper oxide is said to midergo 
reduction. In other words its oxygen is withdrawn. The 
hydrogen used to effect this reduction is called a reducing 
agent. Oxidation and reduction are therefore complemen- 
tary terms. The one consists in adding oxygen to a sub- 
stance, the other in withdrawing it. Reduction as a rule 
is accompanied by oxidation. Thus in the above experi- 
ment, the copper oxide is reduced and the hydrogen is oxi- 
dized. 

In addition to oxygen, hydrogen also combines with a 
number of other elements. Thus it combines with chlorine 
to form hydrochloric acid which is often used as a source 
of hydrogen ; with nitrogen to form common ammonia ; 
with sulphur to form the foul smelling gas, hydrogen sul- 
phide. 

Pure hydrogen produces no injurious results Avhen in- 
haled. Of course one could not live in an atmosphere of it, 
since oxygen is essential to life. 

The Oxyhydrogen Blowpipe. This is a form of apparatus 
used for burning hj-drogen in pure ox^-gen. As previously 
stated, the flame produced b}' the combustion of hydrogen in 



mm. 




Fig. 17 
air is very hot. It is evident that if pure oxygen is substi- 
tuted for air, the temperature reached will be much greater, 
since there will be no inert gases to absorb the heat. The 
oxyhydrogen blowpipe, used to effect this combination, con- 
sists of a small tube placed within a larger one, as shown in 
Fig. ' 17. 



44 Inorganic Chemistry. ^ 

The hydrogen, stored under pressure, generally in steel 
cylinders, is first passed through the outer tube and ignited 
at the end of the tube. The oxygen similarly stored is then 
conducted through the inner tube and mixes with the hydro- 
gen at the end of the tube. In order to produce the maxi- 
mum heat, the hydrogen and oxygen must be present in the 
exact proportion in which they combine, viz., 2 volumes of 
hydrogen to i of oxygen, or by weight, 2 parts of hydrogen 
to 15.8 parts of oxygen. The intensity of the heat may be 
shown by bringing into the flame pieces of metal such as iron 
wire or zinc. These burn with great brilliancy. Even plati- 
num, having a melting point of 1770° may be melted by the 
heat of the flame. 

While the oxyhydrogen flame is intensely hot it is al- 
most non-luminous.' If directed against some infusible sub- 
stance like ordinary lime (calcium oxide) the heat is so in- 
tense that the lime becomes incandescent and glows with a 
brilliant light. This is sometimes used as a source of light 
under the name of "Drummond" or "lime light." 

The Blast Lamp. A similar form of apparatus is com- 
monly used in the laboratory as a source of heat under the 

name "blast lamp" (Fig.- 18). 
This differs from the oxyhydro- 
gen blowpipe, only in the size of 
the tubes. In place of the hydro- 
gen and oxygen the more acces- 
sible coal gas and air are respec- 
" Qdi tively used. The former is com- 
posed largely of a mixture of free 
hydrogen together with gaseous 
compounds of carbon and hydro- 
gen. While the temperature of 
the flame is not so intense as that 
~~~ p. ,g of the oxyhydrogen blowpipe it 

nevertheless suffices fox most 
chemical operations carried out in the laborator3^ 

Uses of Hydrogen. On account of its cost, hydrogen 
is but little used for commercial purposes. It has been 
used, as a material for the inflation of balloons, but gener- 
ally the much cheaper coal gas is now substituted for it. 




Hydrogen. 45 

Even hot air is often used when the duration of ?scension 
is ven short. It has also been largely used as a source of 
heat and light in the oxyhydrogen blowpipe. Where the 
electric current is available, however, this form of apparatus 
has been displaced almost entirely by the electric light and 
electric furnace which are much more economical and more 
powerful sources of light and heat. 



CHAPTER IV. 

COMPOUNDS OF HYDROGEN AND OXYGEN. 
WATER AND HYDl^OGEN DIOXIDE. 

Historical. Water was long regarded as an element. 
In 1 78 1, Cavendish, the discoverer of hydrogen showed that 
it was formed by the union of hydrogen and oxygen. Be- 
ing a believer in the phlogiston theory however he failed 
to interpret his results correctly. A few years later, La- 
voisier repeated Cavendish's experiments and asserted that 
water must be regarded as a compound of hydrogen and 
oxygen. 

The Composition of Water. The composition of 
any substance is ordinarily determined by one of two pro- 
cesses. In the first place it may be possible to separate the 
substance into its constitutents and its composition may thus 
be determined. This process is called analysis. Thus we 
analyze water when we separate it into the elements -which 
compose it. In the second place the substance may be 
formed by the union of its constituent parts. Such a' pro- 
cess is called synthesis. Thus the composition of water may 
be demonstrated by showing that it is the product formed 
by the union of hydrogen and oxygen. 

Analysis and synthesis are therefore opposite processes. The 
one is the separation of a substance into its constituents ; the other 
is the formation of a substance by the union of its constituents. 
These processes may be either qualitative or quantitative. Thus 
qualitative analysis consists simply in ascertaining what elements 
are present : quantitative analysis on the other hand consists in 
determining the exact amount of each element present. Qualitative 
analysis must therefore precede quantitative analysis. 

It has already been shown in the discussion of the 
combustion of hydrogen that water is a compound of hy- 
drogen and oxygen. It remains to ascertain the exact pro- 
portions in- which these elements combine to form water; in 

(46) 



Cojupoiiiids of Hydrogen and Oxygen. 



47 



Q 



other words to determine the quantitative composition of 
water. 

The methods employed in making this determination 
may be classed under the two general heads, (i) Methods 
based on analysis and (2) Methods based on synthesis. 

I. Methods Based on Analysis. As has already been 
indicated, water may be separated into its component 
parts, by means of the electric current. The form of 
apparatus ordinarily used for effecting 
this change is shown in Fig. 19. A plati- 
num wire, to the end of which is attached 
a small piece of platinum foil (about 15 by 
25 mm.) is fused through each of the tubes 
B and D as shown in the figure. 

The stopcocks at the ends of these tubes 
are opened and water, to which has been 
added about one-tenth of its volume of sul- 
B tSi H S phuric acid, is poured into the tube A until 
the side tubes B and D are completely filled. 
The stopcocks are then closed. The plati- 
num wires extending into the tubes B and 
D are now connected with the wires lead- 
ing from two or three di- 
chromate cells joined in se- 
ries. The pieces of platinum 
foil within the tubes thus be- 
come the electrodes and the 
current flows from one to- the other through the acidulated 
water. As soon as the current passes bubbles of gas rise 
from each of the electrodes and collect in the upper part 
of the tubes. The gas rising from the negative electrode 
will be found to have approximately twice the volume of 
that rising from the positive electrode. That the former 
is hydrogen and the latter oxygen may be verified by mak- 
ing the appropriate tests. 




;-#|.|^- 



Fig-. 19 



48 Inorganic Chemistry. 

The process of decomposing a substance by means of the 
electric current is known as electrolysis. The positive elec- 
trode viz., the end of the wire connected with the carbon 
plate of the battery, is called the anode, while the negative 
electrode is called the cathode. As stated above, in the elec- 
trolysis of water the hydrogen escapes from the cathode and 
the oxygen from the anode. Pure water is a non-conductor of 
electricity, but can be made a conductor by the addition of a 
little sulphuric acid. The exact function of the acid will be 
discussed later in the chapter on solutions. 

On close examination, the volume of hydrogen ob- 
tained by the electrolysis of water will be found to be 
slightly more than double that of the oxygen. As has been 
pointed out however oxygen is somewhat more soluble in 
water than hydrogen, and this tends to diminish its volume 
as compared with the volume of hydrogen. It has also been 
found that a small amount of the oxygen liberated is con- 
verted into ozone and hydrogen dioxide. When these facts 
are taken into consideration and appropriate corrections 
made, it can be shown that the volume of the hydrogen ob- 
tained is exactly twice that of the oxygen. 

It is evident that this experiment in itself does not 
prove the composition of water. It simply shows that water 
contains hydrogen and oxygen. It does not prove the ab- 
sence of other elements ; indeed it does not prove the pres- 
ence of hydrogen and oxygen in the proportion in which 
they are liberated. 

That no other elements are present however and that 
water is composed of hydrogen and oxygen in the above 
proportions may be proved in two ways. In the first place 
it may be shown that the weight of the hydrogen and oxy- 
gen obtained by the decomposition of water is exactly equal 
to the weight of water decomposed. In the second place 
the tv/o volumes of hydrogen and the one volume of oxygen 
may be reunited with the formation of water. This leads 
to the discussion of the second of the methods employed for 
the determination of the composition of water. 



Compounds of Hydrogen and Oxygen. 



49 



2. Methods Based on Synthesis. The determination of 
the composition of water by the union of hydrogen and oxy- 
gen is effected in the following way. The combination of 
the gases is brought about in a tube known as a eudio- 
meter (Fig. 20). This is a graduated tube about 60 cm. 
long, 2 cm. wide and closed at one end. Near 
^'^ the closed end two platinum wires are fused 
through the glass, the ends of the wires within 
the tube being separated by a distance of 2 or 
3 mm. The tube is entirely filled with mercury 
and inverted in a vessel of the same liquid. Pure 
hydrogen is passed into the tube until it is about 
■j one-fourth filled. The volume of the gas is then 
I read off on the scale and reduced to standard 
I conditions. About an equal volume of pure oxy- 
B gen is then introduced, the volume again read 

^^^^^ off and reduced to standard conditions. This 
^^^^^H gives the total volume of the two gases and 
Fig. 20 from this the volume of the oxygen intro- 

duced may be determined by subtracting from 
it the , volume of the hydrogen. The combination of 
the two gases is now brought about by connecting the two 
platinum wires with an induction coil and passing a spark 
from one wire to the other. Immediately a slight explosion 
occurs. The mercury in the tube, at first depressed, at 
once rises taking the place of the gases which have 
-combined to form water — the volume of the latter being so 
:small. in comparison with the volume of the gases from 
w^hich it was formed that it may be disregarded. In order 
that the temperature of the residual gas and the liquid may 
become uniform, the apparatus is allowed to stand for a few 
minutes. The volume of the gas is then read off and reduced 
to standard conditions in order that it may be compared with 
those of the hydrogen and oxygen taken. The residual gas 
is then tested in order to ascertain whether it is hydrogen 



50 



Inorganic Chemistry, 



or oxygen. From the data thus obtained the composition 
of the water may be calculated. Thus suppose the data 
obtained were as follows : 

Volume of hydrogen taken 20-3 c.c. 

Volume of hydrogen and oxygen 38.7 c.c. 

Volume of oxygen 18.4 c.c. 

Volume of gas left after combination 

has taken place (oxygen) 8.3 c.c. 

Hence the 20.3 c.c. of hydrogen has combined with 18.4 
— 8.3 or 10. 1 c.c. of oxygen; or approximately 2 volumes 
of hydrogen have combined with i of oxygen. Since oxy- 
gen is 15.88 times as heavy as hydrogen the proportion by 
weight in which the two gases combine is 2 parts of hydro- 
gen to 15.88 of oxygen. 

If the two gases are introduced into the eudiometer in 
the exact proportions in which they combine, after the com- 
bination has taken place the liquid will rise 
and completely fill the tube. Under these 
conditions, however, the tube is very likely 
to be broken by the sudden upward- rush 
of the liquid. Hence in performing the 
experiment care is taken to introduce an 
excess of one of the gases. 

A form of eudiometer (Fig. 21) dif- 
1^ ^ ferent from that given above, is sometimes 

H H used to avoid the calculations necessary to 

^ H H reduce the volumes of the gases to the 

^pJH H same conditions of temperature and pres- 

^^^^^^L^^ sure in order to make comparisons. With 

^^P^ this form it is possible to make the read- 

' ■ ■ ' ings of the volumes under the same con- 

ditions of temperature and pressure, and 
thus compare them directly. The appar- 
atus (Fig. 21) is filled with mercury and 

^. „, the gases introduced into the tube A. The 

Fig. 21 ° 

experiment is carried out as in the preced- 
ing one, except that before taking the reading of the gas vol- 



B 



>^ 



: A 



Compounds of Hydrogen and Oxygen. 



51 



umes mercury is either added to the tube B or withdrawn 
from it by means of the stopcock C until it stands at exactly 
the same height in both tubes. The gas enclosed in tube A 
is then under atmospheric pressure; and since onh* a few 
minutes are required for performing the experiment, the con- 
ditions of temperature and pressure may be regarded as con- 
stant. Hence the volumes of the hydrogen and oxygen and 
of the residual gas may be read off from the tube and directly 
compared. 

By the above method it is possible to determine the pro- 
portion by volume in which hydrogen and oxygen combine. 
From these results, the proportion by weight can be readily 
deduced, knowing the relative weights of the two gases. 
It is possible however to determine directly the proportion 
by weight in which the elements combine. Fig. 22 illustrates 



A=i. 




-^^A. 



Fig. 22 



the apparatus used in making this determination. B repre- 
sents a glass tube containing copper oxide. C and D are 
glass tubes filled with calcium chloride, which has a great 
affinity for water. The tubes B and C including their con- 
tents are carefully weighed, and the apparatus connected as 
shown in the figure. A slow current of pure hydrogen is 
then passed through A, and that part of the tube B which 
contains copper oxide is carefully heated. The hydrogen 
combines with the oxygen present in the copper oxide to 
form water which ns absorbed by the calcium chloride in 
tube C. The operation is continued until an appreciable 



52 Inorganic Chemistry. 

amount of water has been formed. The tubes B and C are 
then again weighed. The loss of weight in the tube B will 
exactly equal the amount of oxygen taken up from the cop- 
per oxide in the formation of the water. The gain in 
weight in the tube C will exactly equal the weight of water 
formed. The difference in these weights will of course 
equal the weight of the hydrogen present in the water 
formed. 

This method for the determination of the composition 
of water was first used by Berzelius in 1820. The work 
was repeated in 1843 by Dumas the average of whose re- 
sults is as follows': 

Weight of water formed 236.36 grains. 

Oxygen given up by the copper oxide 210.04 grains. 

Weight of hydrogen present in water 26.32 grains. 

Hence, according to this experiment, the ratio of hydro- 
gen to oxygen in water is as 26.32 to 210.04 o^ as 2 to 15.96. 

Morley's Results. The most accurate determination 
of the ratio of hydrogen to oxygen in water ever made has 
been recently carried out by the American chemist, Morley. 
The hydrogen and oxygen which combined as well as the 
water formed, were all accurately weighed. Extreme pre- 
cautions were taken to insure pure materials and to eliminate 
all- sources of error. Morley's results show that 2 parts 
of hydrogen by weight combine with 15.88 parts o^ oxygen 
to form water. 

Relation Between the Volume of Aqueous Vapor 
and the Volumes of the Hydrogen and Oxygen Which 
Combine to Form it. 

It was stated above that when hydrogen and oxygen 
are introduced into an eudiometer tube and made to com- 
bine, 'the volume occupied by the resulting water is so 



Compounds of Hydrogen and Oxygen. 



53 




small in comparison with the volumes 
of the gases which combined to form it 
that it could be disregarded. If, how- 
ever the eudiometer tube is surrounded 
by a larger tube (Fig. 23) through 
which is passed the vapor of some 
liquid whose boiling point is above that 
of water, the vapor resulting from the 
union of the hydrogen and oxygen in 
nHJ M the eudiometer tube does not condense 
w^ tsp but remains in the form of steam. It 
^i^^^^ thus becomes possible to compare the 
Fig. 23 volume of steam or water vapor with 

the volumes of the hydrogen and oxygen which combine 
to form it. In this way it has been shown that when two 
volumes of hydrogen and one volume of oxygen combine, 
exactly two volumes of water vapor are formed. It will be 
noted that the relation between these volumes may be ex- 
pressed by whole numbers. The significance of this very 
important fact will be discussed in a subsequent chapter. 

Occurrence of Water. Water not only covers about 
three-fourths of the surface of the earth, and is present in 
the atmosphere in the form of moisture but it is also a 
common constituent of the soil and rocks and almost every 
form of animal and vegetable organism. The human body 
is nearly 70% water. This is derived not only from the 
water which we drink but also from various forms of food 
which contain a large percentage of water. Thus potatoes 
contain about 78% of water, milk 85%, beef over 50%, ap- 
ples 84%, tomatoes 94%. 

Impurities in Water. As has already been shown, 
chemically pure water contains only hydrogen and oxygen. 
Such a water never occurs in nature, however, for being 



54 Inorganic Chemistry. 

a good solvent, it takes up certain substances from the rocks 
and soil with which it comes in contact. When such waters 
are evaporated these substances are deposited in the form 
of a residue. Even rain water which is the purest form oc- 
curring in nature, since it does not come in contact with the 
earth; contains dust particles and gases dissolved from the 
atmosphere. The foreign matter in water is of two kinds 
~viz., mineral, such as common salt and limestone, and or- 
:ganic i. e., the products of animal and vegetable life. 

The amount and nature of the mineral matter present 
in different waters vary greatly, depending on the nature of 
the rocks and soil with which the waters come in contact. The 
more common of the substances present are common salt and com- 
pounds of calcium, magnesium and iron. One liter of the 
average river water contains about 175 mg. of mineral matter. 
Water from deep v/ells naturally contains more mineral mat- 
ter than river water, generally two to three times as much, 
while sea water contains as much as 35,000 mg. to the liter. 

The mineral matter in water does not, save in very ex- 
ceptional cases, render the water injurious to the human 
system. In fact the presence of a certain amount of it is ad- 
vantageous, supplying the mineral matter necessary for the 
formation of the solid tissues of the body. The presence of or- 
ganic matter on the other hand must always be regarded with 
suspicion. Such organic matter may consist not only of the 
products of animal and vegetable life, but also of certain mi- 
croscopic forms of living organisms which are likely to ac- 
company such products. Contagious diseases are known to be 
due to the presence in the body of minute living organisms or 
germs. Each disease is caused by its own particular kind 
of germ. These germs often find their way through the 
sewage from persons afflicted with disease into the water 
supply and it is very often through the drinking water that 
-certain of these diseases, especially typhoid fever, are 
spread. It becomes of great importance therefore to be able 



Compounds of Hydrogen and Oxygen. 



55 



to detect such matter when present in drinking water as 
well as to devise methods whereby it can be removed or at 
least rendered harmless. 

The mineral analysis of a water is, as the name suggests, 
simply the determination of the mineral matter present. San- 
itary analysis, on the other hand, is the determination of the organic 
matter present. The physical properties of a water give no con- 
clusive e^'idence as to its purit}', since a water ma}" be unfit 
for drinking purposes and j^et be perfectly clear and odorless. 
Neither can any reliance be placed on the simple methods often 
given for testing the purity of water. Only the trained chemist 
can carry out such methods of analysis as can be relied upon. 

The Purification of Water. The most effective way 
of purifying natural waters is by the process of distilla- 
tion. This consists in boiling the water and condensing the 
steam. Fig. 24 illustrates the process of distillation as com- 




Fig. 24 

monly conducted in the laboratory. Ordinary water is 
poured into the flask A and boiled. The steam is conducted 
through the condenser B which consists essentially of a 
narrow glass tube sealed within a larger one, the space be- 



56 Inorganic Chemistry. ^ * . 

tween the two being filled with cold water, which is ad- 
mitted at C and escapes at D. The inner tube is thus kept 
cool and the steam in passing through it is condensed. The 
water formed by the condensation of the steam collects in 
the receiver E and is known as distilled water. Such water 
is practically pure, since the impurities being non-volatile,, 
remain in the flask A. 

In preparing distilled water on a large scale, the steam 
is generated in a boiler or other metal container and con- 
densed by passing it through a pipe made of metal, generally 
tin. This pipe is wound into a spiral and surrounded by a 
current of cold water. Distilled water is' used by the chemist 
in almost all of his work. It is also used in the manufacture 
of artificial ice and for drinking water. 

In preparing distilled water, it is evident that if the nat- 
ural water contains some substance which is volatile, its vapor 
will pass over and be condensed with the steam, so that 
the distillate will not be pure water. Even such mixtures,, 
however, may generally be separated by repeated distillation. 
Thus if a mixture of water (boiling point ioo°) and alcohol 
(boiling point 78°) is distilled, the alcohol, having the lower 
boiling point, ten-ds to distill first, followed by the water. The 
separation of the two is not perfect, however, but may be 
made nearly so by repeated distillations under certain conditions 
that can not properly be discussed in an elementary treatise. 
The process of separating a mixture of volatile substances by 
distillation is known as fractional distillation. 

The process of distillation practically removes all non- 
volatile foreign matter, mineral as well as organic. In puri- 
fying water for drinking purposes however it is only neces- 
sary to eliminate the latter or to render it harmless. This is 
ordinarily done by either one of two different processes, viz., 
filtration and boiling. In filtration the water is passed through 
some medium which will retain the organic matter. Ordi- 
nary charcoal is a porous substance and will condense within 
its pores the organic matter in water if brought in contact 
with it; hence its use in the construction of filters. 



Coinpotinds of Hydrogen and Oxygen. 57 

Such niters in order to be effective must be kept clean, 
since it is evident that the charcoal is useless after its pores 
are filled. A more effective type of filter is the Chamberlain- 
Pasteur. In this, the water is forced through a porous, 
cylindrical cup, the pores being so minute as to strain out 
the organic matter. 

A simpler and equally effective method for purifying 
water for drinking purposes consists in boiling the water. 
It is the germs in water that render it dangerous to health. 
These germs are living forms of matter. If the water is 
boiled, the germs are killed and the water rendered safe. 

While these germs are destro3'ed b}^ heat, cold has little 
effect upon theni. Thus Dewar, in working with liquid h3^dro- 
gen, exposed some of these minute forms of life to the tempera- 
ture of boiling hydrogen ( — 252°) without killing them. It is 
clear from this that ice frozen from impure water can not 
be regarded as free from impurities. 

The above processes, viz., filtration and boiling, remove 
at most only a small amount of the mineral matter dissolved 
in the water. The main mineral ingredient removed is lime- 
stone. 

Reference was made under the discussion of the uses of 
oxygen, to the self-purification of running water, due to the 
oxidation of the organic matter present through the agency 
of the atmospheric ox3'gen. While water is purified in this 
way, j^et the method can not be relied upon to purify a con- 
taminated water so as to render it safe for drinking purposes. 

Water which has percolated through the soil to any dis- 
tance is purified by the absorbent action of the soil. Advan- 
tage is taken of this fact in purifying the water supply of cities. 
Large filtration beds are prepared, from sand and gravel and 
the water is allowed to flow through these. Cities supplied with 
water purified in this way are practically free from typhoid 
fever. 

Physical Properties. Pure w^ater is an odorless and 
tasteless liquid, colorless in thin layers but having a bluish 
ting£ when observed through a considerable thickness. Un- 



58 • Inorganic Chemistry. 

der the normal pressure of one atmosphere, water boils at 
ioo° and solidifies at o°. If the pressure is increased the 
boiling point is raised and the freezing, point is lowered. 
When water is cooled it steadily contracts until a tempera- 
ture of 4° is reached ; it then expands. Hence ice is lighter 
than water and floats upon it. 

Chemical Properties. It will be recalled that the 
imion of hydrogen with oxygen to form water is attended 
by the liberation of a large amount of heat. Such a 
body is very difficult to decompose by heat alone since ex- 
periment has shown that it requires as much heat to effect 
the decomposition of a substance as is given off when the 
substance is formed. In the case of water, this decomposi- 
tion begins near 1000° but is only about half complete at 
2500°. If the heat is withdrawn combination of the ele- 
ments of course takes place again. While water is thus very 
stable towards heat it may be decomposed as already pointed 
out, by the electric current and by the action of certain ele- 
ments as sodium, iron and carbon. 

Many substances such as sugar or salt, when brought 
into contact with water are taken up or dissolved by it, a 
solution of the substance in water being formed. The na- 
ture of solutions will be discussed later in the chapter de- 
voted to this subject. It is only necessary to add here that 
compounds in solution undergo chemical changes much 
more readily than when in the dry or solid forms. Even the 
slightest trace of moisture, in some unexplained way, often 
materially assists in causing chemical changes to take place. 
Thus phosphorus usually burns in oxygen with great en- 
ergy ; but if extreme precautions are taken to remove all 
traces of moisture, the two elements will not combine at all. 
Again the kindling temperature of hydrogen and oxygen 
tinker ordinary conditions is about 600° ; but if the gases 



Compounds of Hydrogen and Oxygen. 59 

are perfectly dry the kindling temperature is considerably 
raised. Many other examples of this kind might be given. 
Water of Crystallization. When solutions of solid 
■compounds are allowed to stand, the water evaporates and 
the solid compound separates out, often in the form of crys- 
lals. It has been found that the crystals of many com- 
pounds, although perfectly dry, give up a definite amount of 
water when heated, the substance at the same time losing 
its crystalline form. Such water is called water of crystal- 
lization. This varies in amount with different compounds 
but is perfectly definite for the same compound. Thus if a 
crystal of copper sulphate is strongly heated in a tube, water 
is evolved and condenses on the sides of the tube, the crys- 
tal crumbling to a light powder. The weight of the water 
evolved is always equal to exactly 36.07 % of the weight of 
copper sulphate crystals heated. The water must there- 
fore be in chemical combination with the substance compos- 
ing the crystal ; for if simply mixed with it or adhering to it, 
not only would the substance appear moist but the amount 
present , would undoubtedly vary. The combination how- 
ever must be a very weak one since the water is often ex- 
pelled by even a gejitle heat. Indeed in some cases, the water 
is evolved on simple exposure to air. Such compounds are 
said to be efflorescent. Thus a crystal of ordinary sodium 
sulphate (Glauber's salt) on exposure to air crumbles to a 
iine powder, owing to the escape of its water of crystalliza- 
tion. Other substances have just the opposite property, viz., 
they absorb moisture when exposed to the air. Thus if a 
bit of dry calcium chloride is exposed to the air, in the 
course of a few hours it will have absorbed sufficient mois- 
ture to dissolve it. Such substances are said to be deliques- 
cent. A deliquescent body serves as a good drying or 
desiccating agent. Thus we have already employed calcium 



6o Inorganic Chemistry. 

chloride as an agent for absorbing the moisture from 
hydrogen. 

Water of crystallization must be carefully distinguished 
from water which is mechanically inclosed in a crystal and 
which can be. removed by powdering the crystal and drying. 
Thus when crystals of common salt are heated, the water 
inclosed in the crystal is changed into steam and bursts the 
crystal with a crackling sound. Such crystals are said to 
decrepitate. That this water is not combined is proven by 
the fact that the amount of it present varies and that it has 
all the properties of water. 

Uses of Water. The importance of water in its re- 
lation to life and commerce is too well known to require 
comment. Its importance to the chemist has also been 
pointed out. It simply remains to call attention to the fact 
that it is used as a standard in many physical measurements. 
Thus o° and ioo° on the centigrade scale are respectively 
the freezing and the boiling points of water under normal 
pressure. The weight of i c.c. of water at its point of great- 
est density is the unit of weight in the metric system viz., the 
gram. It is also taken as the unit for the determination of 
the density of liquids and solids as well as for the measure- 
ment of amounts of heat. 

HYDROGEN DIOXIDE. 

As has been shown, two parts by weight of hydrogen 
combine with 15.88 parts by weight of oxygen to form 
water. It is possible however to obtain a second compound 
of hydrogen and oxygen differing from water in composi- 
tion in that two parts by weight of hydrogen are com- 
bined with 2 X 15-88 or 31.76 parts of oxygen. This com- 
pound is called hydrogen dioxide or hydrogen peroxide the 
prefixes "di" and ''per" signifying that it contains more 
oxygen than hydrogen oxide which is the chemical name for 
water. 



Compounds of Hydrogen and Oxygen. 6i 

Preparation. Hydrogen dioxide can not be prepared 
economically by the direct union of hydrogen and oxygen; 
indirect methods must therefore be used. It is commonly 
prepared by the action of sulphuric acid on barium dioxide. 
The change which takes place may be indicated as follows : 

f hvdrogen f , . f barium f' , , 

' . barmm I , , nyaro2:en 

^ sulphur _-<^ =-<^ sulphur — < 

1 ^ 1 oxv2:en ^, "■ ' \ oxvgen 

I oxygen l I o^y§"^^^ L 

(sulphuric (barium (barium (hydrogen 

acid) dioxide) sulphate) dioxide) 

In other words the barium and hydrogen in the two 
compounds exchange places. By this method a dilute solu- 
tion of the dioxide in water is obtained. Inasmuch as the 
boiling point of the dioxide is higher than that of water it is 
possible to separate the two by fractional distillation. This 
is attended with great difficulties however, since the pure 
dioxide is explosive. The distillation is carried on under 
diminished pressure so as to lower the boiling points as 
much as possible ; otherwise the high temperature will de- 
compose the dioxide. 

Properties and Uses. The pure hydrogen dioxide 
is a colorless syrupy liquid having a density of 1.49. Its 
most characteristic property is the ease with which it decom- 
poses into water and oxygen. Two parts by weight of hy- 
drogen are capable of holding firmly only 15.88 parts of 
oxygen. The extra 15.88 parts present in hydrogen diox- 
ide are therefore easily evolved, the compound breaking 
down into water and oxygen. This decomposition is at- 
tended by the evolution of considerable heat. In dilute so- 
lutions hydrogen dioxide is fairly stable, although such solu- 
tions should be kept in a dark cool place since both heat and 
light aid in its decomposition. These solutions are used 
largely as oxidizing agents i. e., they readily give up oxygen 
to other substances which have an affinity for it. A 3% 



62 Inorganic Chemistry. 

solution in water is commonly used in medicine as an anti- 
septic. When brought in contact with organic matter, it 
decomposes, the evolved oxygen destroying the germs of 
disease. 

It is a noteworthy fact that the decomposition of hydro- 
gen dioxide may be brought about by the presence of cer- 
tain metals. Thus if a piece of platinum is introduced into a 
concentrated solution of the dioxide the decomposition takes 
place with explosive violence. Just how the metal efifects this 
change is not understood, since it undergoes no apparent change. 



CHAPTER V. 
THE ATOMIC THEORY. 

Fundamental Laws of Matter. 

1. Lazv of Conserz'ation of Matter. In the introduc- 
tory chapter reference was made to the law of conservation 
of matter, which states that in all the changes which sub- 
stances can be made to undergo, no matter is either created 
or destroyed. Attention must now be directed to two 
equally important laws relating to the composition of chem- 
ical compounds. 

2. Law of Definite Composition. In the earlier days 
of chemistry there was much discussion as to whether the 
composition of a given compound is always precisely the 
same, or whether it is subject to some variation. Two 
Frenchmen, Berthollet and Proust, were the leaders in this 
discussion, and a great deal of most useful experimenting 
was done to decide the question. Their experiments, as 
well as all succeeding ones, have shown that the composition 
of a pure chemical compound is always exactly the same. 
Water obtained by melting pure ice, condensing steam, burn- 
ing hydrogen in oxygen, has always 11.19% hydrogen, and 
88.81% oxygen in it. Red oxide of mercury from whatever 
source it is obtained, contains 92.6% mercury and 7.4% oxy- 
gen. This truth may be summed up in the statement that 
the composition of a chemical compound never varies. This 
statement is called the law of definite proportion. 

3. Law of Multiple Proportion. It has already been 
noted however, that hydrogen and oxygen combine in two 
different proportions to form water and hydrogen dioxide. 
It will be observed that this fact does not contradict the law 
of definite proportions, for entirely diiTferent substances are 

(63) 



64 Inorganic Chemistry. 

formed. These compounds differ from each other in com- 
position, but the composition of each one is always con- 
stant. This abihty of two elements to unite in more than 
one proportion is very frequently observed. Carbon and 
oxygen combine in tw^o different proportions ; nitrogen and 
oxygen combine to form as many as five distinct compounds 
each with its own precise composition. 

In the first decade of the last century, John Dalton, an 
English school teacher and philosopher, endeavored to find 
some rule which holds between the ratios in which two 
given substances combine. His studies brought to light a 
very simple relation which the following examples will make 
clear. In water, as has been already stated, the hydrogen and 
oxygen are combined in the ratio of 2 parts by weight of 
hydrogen to 15.88 parts by weight of oxygen. In hydrogen 
dioxide the 2 parts by weight of hydrogen are combined with 
31.76 parts by weight of oxygen. The ratio between the 
amounts of oxygen which combine with the same amount of 
hydrogen to form respectively water and hydrogen dioxide 
is therefore 15.88: 31.76, or I : 2. 

Similarly, the element iron combines with oxygen to 
form two oxides, one of which is black and the other red. By 
analysis it has been shown that the former contains 2 parts 
by weight of oxygen combined with 7 parts by weight of 
iron, while the latter contains 3 parts by weight of oxygen 
combined with 7 parts by weight of iron. Here again we 
find that the amounts of oxygen which combine w^ith the 
same amount of iron to form the two compounds are in 
the ratio of small whole numbers, viz. 2 : 3. 

Many other examples of this simple relation might be 
given, since it has been found to hold true in all cases 
where more than one compound is formed from the same ele- 
ments. Dalton's law of multiple proportion states these facts 
as follows : When anv two elements, A and B are able to 



The Atomic Theory. 65 

form more than one compound, the amounts of B which 
unite with any fixed amount of A bear the ratio of small, 
w^hole numbers to each other. 

The Atomic Theory. These three generalizations 
are called lazvs, because they express in concise language 
truths which are found by careful experiment to hold good 
in all cases. They do not offer any explanation of the facts, 
b)u'- merely state them. The human mind, however, does 
not rest content with the mere bare facts, but seeks ever 
to learn the explanation of the facts. A suggestion which 
is offered to explain such a set of facts is called an hy- 
pothesis. The suggestion which Dalton offered to explain 
the three laws of matter was prompted by his view of the 
constitution of matter, and it involves three distinct assump- 
tions in regard to the nature of matter and chemical action. 
Dalton could not prove these assumptions to be true, but he 
saw that if they were true, the laws of matter become very 
«asy to understand. 

Three Assumptions upon v^hich Dalton's Hypothe- 
sis Rests. The three assumptions which are involved in 
.Dalton's atomic hypothesis are these : 

1. AlLelemerits consist of minute particles which Dal- 
ton called atoms. According to this view matter is not con- 
tinuous, but is made up of minute units. 

2. All atoms of the same element have equal masses ; 
those of different elements have different masses. In any 
<:hange to which an atom is subjected, its mass does not 
•char e. 

3. When chemical union takes place between two ele- 
ments, the action consists in a definite number of atoms of 
the one element uniting with a definite number of the atoms 
of the other element to form a small particle of the com- 
pound. 



66 Inorganic Chemistry. 

Dalton called the particles of the compound, as well as 
those of the elements, atoms. Later Avogadro, an Italian 
scientist, pointed out the fact that the two are different,, 
since the smallest particle of an element is a unit, while that 
of a compound must have at least two units in it. He sug- 
gested the name molecule for the least particle of a com- 
pound which can exist, retaining the name atom for the 
smallest particle of an element. 

Supposing these three assumptions to be true, let us 
now see how they serve to explain the laws of matter. 

1. Explanation of the Lazu of Conservation of 
Mass. It is evident that if the atoms never change their 
masses in any change which they undergo, the total mass 
of matter can never change and the law of conservation 
of mass must follow. 

2. Explanation of the Lazv of Definite Proportion. 
According to the third supposition, when iron com- 
bines with sulphur the union is between definite num- 
bers of the two kinds of atoms. In the simplest case one 
atom of the one element combines with one atom of the 
other. If the sulphur and the iron atoms never change their 
respective masses when they unite to form a molecule of 
iron sulphide, all iron sulphide molecules will have equal 
amounts of iron in them, and also of sulphur. Conse- 
quently any mass made up of iron sulphide molecules, will 
have the same fraction of its weight iron as do the indi- 
vidual iron sulphide molecules. Iron sulphide, from what- 
ever source will have the same composition, which is the law 
of definite proportion. 

3. Explanation of the Law of Mnltiple Proportion. 
But this simplest case may not always be the only one. 
Under other conditions one atom of iron might combine with 
twp of sulphur to form a molecule of a second compound. 
In such a case the one atom of iron would be in combina- 



The Atomic Theory. 6y 

tion with twice the mass of sulphur that is in the first com- 
pound, since the sulphur atoms all have equal masses. 
What, is true for one molecule will be true for any number 
of them ; consequently when definite quantities of these two 
compounds are found to contain the same amount of iron, 
one will have twice as much sulphur as the other. 

The combination between the atoms might occur in 
other simple ratios. Thus two atoms of the one element 
might combine with three of the other or two with five. In 
any case the union would produce compounds having defi- 
nite composition, and on selecting such' numbers of mole- 
cules of the two compounds as would contain equal numbers 
of the atoms of one constituent, the numbers of the atoms 
of the other constituent would have to bear simple ratios to 
each other. And if the numbers of the atoms are in simple 
ratio to each other, the masses of them would have to be,, 
since the weights of the atoms of a given element are all 
equal. Consequently if we select such amounts of the two 
compounds as will contain equal masses of one of the ele- 
ments, the masses of the other element in the two com- 
pounds w^ill have to bear a simple ratio to each other. 

Testing the hypothesis. Efforts have been made 
to find compounds which do not conform to these laws, but 
all such attempts have resulted in failure. If such com- 
pounds were to be found, the laws would be no longer true, 
and the hypothesis of Dalton would cease to possess value. 
When an hypothesis has been tested in every way that ex- 
periment can test it, and is still found to be in harmony 
with the facts in the case it is termed a theory. We now 
speak of the atomic theory rather than of the atomic hy- 
pothesis. 

Value of a Theory. The value of a theory is two- 
fold. , It aids in the clear understanding of the laws of 



68 Inorganic Chemistry. 

nature because it gives an intelligent idea as to why these 
laws should be in operation. 

A theory also adds much to our knowledge. It usu- 
ally happens that in testing a theory much valuable work is 
done, and many new facts are discovered. Almost any 
theory in explaining given laws, will involve a number of 
consequences apart from the, law it seeks to explain. Ex- 
periment will soon show whether these facts are as the 
theory predicts they will be. Thus Dalton's atomic theory 
predicted many properties of gases which experiment has 
since verified. 

Atomic Weights. It v^ould be of great advantage 
in the study of. chemistry if we could determine the weights 
.of the different kinds of atoms. That this cannot be done 
directly is evident. They are so small that they cannot be 
seen even with a most powerful microscope. No balance 
can weigh such minute objects. It is possible, however, 
to determine their relative weights — that is how much 
heavier one is than another. These relative weights of the 
atoms are spoken of as the atomic weights of the elements. 

If elements were able to combine in only one way — 
one atom of one with one atom of another — the prob- 
lem of determining the atomic weights would be 
very simple. We should merely have to take some one 
convenient element as a standard, and find by experi- 
ment how much of each other element would combine with 
a fixed amount of it. The ratios thus formed would be the 
same ratios as those between the atoms of the elements, and 
thus we should have their relative atomic wei^g-hts. The 
law of multiple proportion calls attention to the fact that the 
atoms combine in other ratios than i :i, and there is no direct 
way of telling which one of the several compounds in a 
given case is the one consisting of a single atom of each 



The Atomic Theory. ^ 69 

element. If some way were to be found of telling how 
much heavier the entire molecule of a compound is than the 
atom chosen as standard, that is of determining the molecular 
weights of compounds, the problem could be solved, though 
its solution would not be an entirely simple matter. Tliere 
are ways of obtaining this information, that is of determin- 
ing the relative molecular weights of compounds, and there 
are other experiments which throw light upon the relative 
weights of the atoms directly. These methods cannot be 
described until the facts upon which they rest have been 
studied. It will be sufficient for the present to assume that 
these methods are trustworthy. 

Standard for Atomic Weights. The figures follow- 
ing the symbols of the elements in the table on the front 
cover of the book give the relative weights of the atoms of 
the elements. They are not weights in grams or any other 
system of units. They merely mean that if we were to 
assign the arbitrary value of one unit of any kind to the atom 
of hydrogen, we should have to give the value 55.6 to the 
atom of iron, 31.83 to the atom of sulphur, etc. Whatever 
may be the actual weight of the atom of hydrogen and of 
the atom of iron, the ratio of their weights is i : 55.6. 

Since the atomic weights are merely relative to some 
one element chosen as a standard, it is evident that any one 
of the elements may serve as this standard, and that any 
convenient value may be assigned to its atom. At one time 
oxygen was taken as the standard with the value 100 and the 
other elements were given weights to agree with this. It 
would seem to be more rational to take the element of smallest 
atomic weight as the standard, and give it unit value; ac- 
cordingly hydrogen is assumed as the standard, with an 
atomic weight of i. On this basis oxygen is 15.88. Fof 
many reasons oxygen serves better as a standard, and if the 
value .t6 is given to it, hydrogen becomes 1.008. There is 



70 Inorganic Chemistry. 

little difference in the two sets of values which will be ob- 
tained from these two standards, and since the reason for 
assuming hydrogen as the unit is a little easier to understand, 
we shall keep it as the standard in the chapters which follow. 



CHAPTER VI. 
CHEMICAL EQUATIONS AND CALCULATIONS. 

Formulas. Since the molecule of any chemical com- 
pound consists of a definite number of atoms, and this num- 
ber never changes without destroying the identity of the 
compound, it is very convenient to represent the composi- 
tion of a compound by indicating the composition of its 
molecules. This can easily be done by using the symbols 
of the atoms to indicate the number and the kind of the 
atoms which constitute the molecule. HgG will in this way 
represent mercuric oxide, a molecule of which has been 
found to contain one atom each of mercury and oxygen. 
H2O will represent water, the molecules of which consist 
of one atom of oxygen and two of hydrogen, the subscript 
figure indicating the number of the atoms of the element 
whose symbol precedes it. HoSO^ will stand for sulphuric 
acid, the molecules of which contain two atoms of hydrogen, 
■one of sulphur and four of oxygen. The combination of 
symbols which represents the molecule of a substance is 
called its formula. 

Equations. When a given substance undergoes a 
chemical change, it is possible to represent this change by the 
use of such symbols and formulas. It has already been 
seen that mercuric oxide decomposes when heated to form 
mercury and oxygen. This may be expressed very briefly 
in the form of an equation : 

(1) HgO = Hg + 0. 

When water is electrolyzed, two new substances, hy- 
drogen and oxygen, are formed from it. This statement in 
the form of an equation is : 

(2) H,0 = 2H + O. 

(71) 



^2 ^ Inorganic Chemistry. 

The coefficient before the symbol for hydrogen indi- 
cates that a single molecule of water yields two atoms of 
hydrogen on decomposition. In like manner the combina- 
tion of sulphur with iron is expressed by the equation : 

(3) Fe + S = FeS. 

The decomposition of potassium chlorate by heat takes 
place as represented by the equation : 

(4) KCIO3 = KCl + 3 O. 

Reading of Equations. Since equations are simply 
a kind of short hand way of indicating chemical changes 
which occur under certain conditions, in reading an equation 
the full statement for which it stands should be given. 
Equation (i) should be read "mercuric oxide when heated 
gives mercury and oxygen." Equation (2) is equivalent 
to the statement 'Svhen electrolysed, water produces hydro- 
gen and oxygen." Equation (3), "when heated together, 
iron and sulphur unite to form iron sulphide ;" equation 
(4), ''potassium chlorate when heated yields pota'ssium 
chloride and oxygen." 

How Equations are Derived. In order to write such 
equations correctly a considerable amount of exact knowl- 
edge is required. Thus in equation (i) the fact that 
red oxide of mercury has the composition represented by 
the formula HgO, that it is decomposed by heat, that in this 
decomposition mercury and oxygen are formed and no other 
products — all these facts must be ascertained by exact ex- 
periment before the equation can be written. An equation 
expressing these facts will then have much value. 

Having obtained an equation describing the conduct of 
mercuric oxide on being heated, it will not do to assume 
that similar oxides will behave in like manner. Iron oxide, 
FeCT, resembles mercuric oxide in many respects, but it 



Chemical Equations and Calculations. 73 

undergoes no change at all when heated. Manganese di- 
oxide, the black substance used in the preparation of oxygen, 
has the formula MnOs- When this substance is heated, 
oxygen is set free, but none of the metal manganese is 
formed. Instead a different oxide of manganese containing 
less oxygen is produced. The equation representing the re- 
action is : 

'sMnO, = MngO, + 2O. 

Substitution. The equations so far considered, rep- 
resent the simple union of two elements, or the decomposi- 
tion of a compound into other substances. It is sometimes 
possible for an element in the free state to act in such a way 
upon a compound that it takes the place of one of the ele- 
ments of the compound, liberating it in turn. In the -study 
of the element hydrogen it was pointed out that hydrogen 
is most easily prepared by the action of .sulphuric or hydro- 
chloric acid upon zinc. AA^hen sulphuric acid is used, a sub- 
stance called zinc sulphate, having the composition repre- 
sented by the formula ZnSO^ is formed, together with hy- 
drogen. The equation is : 

(5) Zn + H,SO, ^ ZnSO, + 2H. 

AA'hen hydrochloric acid is used, zinc chloride and hy- 
drogen are the products of reaction. 

(6) Zn + 2HCI = ZnCL + 2H. 

These two reactions are quite similar as is apparent 
from an examination of the two equations. In both cases 
an atom of zinc replaces two atoms of hydrogen in the acid, 
and the hydrogen escapes as a gas. When an element in 
the free state, such as zinc, takes the place of some one ele- 
ment in a compound, setting it free from chemical combina- 
tion, the act is called substitution. 



74 Inorganic Chemistry. 

Other reactions illustrating substitution are the action 
of sodium on water, 

(7) Na + H,0 = NaOH + H 

and the reaction of water upon heated iron: 

(8) 3Fe + 4H2O = FegO^ + 8H. 

Double Decomposition. When barium dioxide, 
which has the composition represented by the formula BaOa, 
is treated with sulphuric acid, two compounds are formed, 
viz., hydrogen dioxide, HoO^, and barium sulphate, BaSOi- 
The equation is 

(9) BaO, + H.SO4 ^ BaSO, + H^O,. 

In this reaction it will be seen that the two elements 
"barium and hydrogen simply exchange places. Such a re- 
action is called a double decomposition or metathesis. We 
shall meet with many examples of this kind of chemical 
reactions. 

These four kinds of chemical reactions, viz., direct com- 
b)ination, decomposition, substitution and double decompo- 
sition, are the most common types of chemical reactions. 

Chemical Equations are Quantitative. The use of 
symbols and formulas in expressing chemical changes has 
another great advantage. Thus according to the equation, 
H^O ^ 2 H + O, I molecule of water is decomposed into 
2 atoms of hydrogen and i atom of oxygen. But as al- 
ready stated, the relative weights of the atoms are known, 
that of hydrogen being i while that of oxygen is 15.88. The 
molecule of water being composed of 2 atoms of hydrogen 
and I atom of oxygen must therefore weigh relatively 2 -j- 
15.88 or 17.88. The amount of hydrogen in this molecule 
must therefore be -^^j^ or 11.18% of the whole while 
the amount of oxygen must be -yj-^ or 88.82% of the 
whole. Now since any amount of water is simply the sum 



Chemical Equations and Calculations. 75 

of a certain number of molecules of water, it is plain that 
the numbers representing the relative amounts of hydrogen 
and oxygen present in a molecule must likewise express the 
relative amounts of hydrogen and oxygen present in any 
quantity of water. Thus for example, in 20 grams of water 
there are yf^ X 20 or 2.236 grams of hydrogen and 
^^ X 20 or 17.764 grams of oxygen. It will be ob- 
served that these results in reference to the composition of 
water express the facts obtained by direct experiment as 
given in the chapter on water. 

It may be easier for some students to make the above 
calculation by proportion. Since the molecule of water and 
the two atoms of hydrogen which it contains have the ratio 
by weight of 17.88 : 2, any mass of water has the same ratio 
between its total weight and the weight of the hydrogen 
in it. Hence to find the number of grams (x) of hydrogen 
in 20 grams of water, we have the proportion 

17.88 : 2 : : 20 grams :x (grams of hydrogen). 

Solving for x we get 2.236 for the number of grams 
of hydrogen. Similarly to find the amount (x) of oxygen 
present in the 20 grams of water we have the proportion 

17.88: 15.88: :2o:x 

from which we find that x = 17.764 grams. 

Again, suppose we wish to find how much oxygen can 
be obtained from 15 grams of mercuric oxide. The equation 
representing the decomposition of mercuric oxide is HgO= 
Hg+O. The relative weights of the mercury and oxygen 
atoms are in round numbers respectively 199 and 16. The rel- 
ative weight of the mercuric oxide molecule must therefore 
be the sum of these, or 215. The molecule of mercuric oxide 
and the atom of mercury which it contains have the ratio 
215:16. This same ratio must therefore hold between the weight 
of any given quantity of mercuric oxide and the oxygen which 
it contains. Hence to find the weight of oxygen in 15 grams of 



76 Inorganic Chemistry. 

mercuric oxide we have the proportion 215 : 16 :: 15 : x (grams 
of oxygen). 

On the other hand suppose we wish to prepare say 20 
grams of oxygen. The problem is to find out what quantity 
of mercuric oxide will yield 20 grams of oxygen. The fol- 
lowing proportion evidently holds good: 

215 : 16 : : x (grams of mercuric oxide) :20, 

from which we get x = 268.7. 

In the preparation of hydrogen by the action of sulphuric 
acid upon zinc, according to the equation 

Zn+H2S04=ZnS04+2H, 

suppose that 50 grams of zinc are available; let it be required 

to calculate the weight of hydrogen which can be obtained. 

It will be seen that one atom of zinc will liberate two atoms 

of hydrogen. The ratio by weight of a zinc to an hydrogen 

atom is 65 : I ; of a zinc atom to two hydrogen atoms, 65 : 2. 

Zinc and hydrogen will be related in this reaction in this same 

ratio, however many atoms of zinc are concerned. Consequently 

in the proportion 

65 : 2 : : 50 : X, 

x will be the amount of hydrogen set free by 50 grams of 
zinc. The amount of zinc sulphate produced at the same' time 
can be found from the proportion 65 : 161 ::5o:x, where 161 
is the molecular weight of the zinc sulphate, and x the weight 
of zinc sulphate formed. In like manner the amount of sul- 
phuric acid used up can be calculated from the proportion 

65 : 98 : : 50 : X. 

These simple calculations are possible because the sym- 
bols and formulas in the equations represent the relative 
quantities of * substances concerned in a chemical reaction. 
When once the relative weights of the atoms have been de- 
termined, and it has been agreed to allow the symbols to 
stand for these relative weights, an equation or formula 
making use of the symbols becomes a statement of a definite 
numerical fact, and calculations can be based ori it. 

Chemical Equations Not Algebraic. Although 
chemical equations are quantitative, it must be clearly 



Chemical Equations and Calculations. yy 

understood that they are not algebraic. A glance at 
the equations 7-|-4= ii»8 + 5^9 + 4 will show at once 
thatjhey are true. The equations HgO = Hg + O, FeO = 
Fe + O are equally true in an algebraic sense, but experi- 
ment shows that only the first is true chemically. Only 
such equations as have been found by careful experiment to 
express a real chemical transformation, true both for the 
kinds of substances as well as for the amounts, have any 
value. 

Chemical formulas and equations therefore, are a con- 
cise way of representing qualitatively and quantitatively 
facts which have been found by experiment to be true in 
reference to the composition of substances and the changes 
which they undergo. 

Heat of Reaction. Attention has frequently been 
directed to the fact that chemical changes are usually ac- 
companied by heat changes. In general it has been found 
that in every chemical action heat is either absorbed or 
given off. The equations so far employed take no notice 
of these heat changes ; and in order to record them it is 
first necessary to find a way to measure heat accurately. 
This cannot be done by the use of a thermometer alone, 
since the thermometer measures the intensity of heat, not 
its quantity. 

The easiest way to measure a quantity of heat is to note 
how warm it will make a definite amount of a given sub- 
stance chosen as a standard. Watfer has been chosen as the 
standard and the unit of heat, called a calorie, is defined as 
the amount of heat required to raise the temperature of i 
gram of water from o° to i°. 

By means of this unit it is easy to indicate the heat 
changes in a given chemical reaction. The reaction 

2 H + O = H,0 + 68300 cal, 



78 Inorganic Chemistry. 

means that when 2 grams of hydrogen combines with 15.88 
grams of oxygen, to form 17.88 grams of water, 68300 
units of heat are set free. 

C + 2S = CS2 — 19000 cal. 

means that an expenditure of 19000 heat units is required 
to cause 12 grams of carbon to unite with 64 grams of sul- 
phur, to form yG grams of carbon disulphide. 

Conditions of a Chemical Action not Indicated by 
Equations. Equations do not tell the conditions under 
which a reaction will take place. The equation 

HgO = Hg + 

does not tell us that it is necessary to Jceep the mer- 
curic oxide at a high temperature that the decomposition 
may go on. The equation 

Zn + 2 HCl = ZnClg + 2 H 

in no way indicates the fact that the hydrochloric acid must 

be dissolved in water before it will act upon the zinc. From 

the equation - 

H 4- CI = HCl 

it would not be suspected that the two gases hydrogen and 
chlorine will unite instantly in the sunlight, but will stand 
mixed in the dark a long time without change. It will there- 
fore be necessary to pay much attention to the details of the 
conditions under which a given reaction occurs, as well as 
to the expression of the reaction in the form of an equation. 



CHAPTER VII. 

NITROGEN AND THE RARE ELEMENTS, ARGON -> 
NEON ^ HELIUM — KRYPTON — XENON. 

HistoricaL Nitrogen was discovered by the Eng- 
lish chemist Rutherford in 1772. A little later Scheele 
showed it to be a constituent of air and Lavoisier gave 
it the name azote, signifying that it would not support 
life. The name nitrogen was afterwards given it be- 
cause of its presence in saltpetre or nitre. The term 
azote and symbol Az are still retained by the French 
chemists. 

Occurrence. Attention has already been called to 
the fact that the air is composed principally of oxygen and 
nitrogen in a free state, about 78 parts by volume out of 
every 100 parts being nitrogen. Nitrogen also occurs in 
nature in the form of potassium nitrate, KNO3, (commonly 
called saltpetre or nitre) as well as in sodium nitrate, 
NaNOg. Nitrogen is also a necessary constituent of living 
forms of matter. 

Methods of Preparation. 

I. Preparation from Air. Usual laboratory method. 
Nitrogen is commonly prepared in the laboratory by remov- 
ing' the oxygen from a confined portion of air. This may 
be done by burning some substance in the confined air which 
has a very strong affinity for oxygen. Such a substance 
must be chosen however as will combine with the oxygen 
to form a product which is not a gas : otherwise the nitro- 
gen present will be mixed with the gaseous products of 
combustion. The substance ordinarily used is phosphorus. 
This is placed in a little porcelain dish attached to a cork 

(79) 



8o 



Inorganic Chemistry. 




Fig. 25 



and floated on water. (Fig. 25.) 
It is then ignited by contact with 
a hot wire and immediately a bell 
jar or bottle is brought over it so 
as to confine a portion of the air. 
The phosphorus combines with the 
oxygen to form an oxide of phos- 
~ phorus known as phosphorus pent- 
oxide. This is a white solid which 
floats about in the bell jar but on standing is all absorbed 
by the wafer, leaving the nitrogen. The withdrawal of the 
oxygen is indicated by the rising of the water in the bell jar. 
The oxygen present in the air may also be removed by 
passing air slowly through a heated tube containing copper. 
The copper combines with the oxygen to form copper oxide. 
Inasmuch as air, in addition to oxygen and nit- 
rogen contains small amounts of other gases, and since 
the phosphorus as well as the copper removes only 
the oxygen, it is evident that the nitrogen ob.tained 
by this method is never quite pure. About 1% of 
the product is composed of other gases from which 
it is very difficult to separate the nitrogen. The nitrogen 
so obtained however may be used for a study of some of its 
properties since these are not materially affected by the 
presence of the other gases. 

2. Preparation from Compounds of Nitrogen. Pure 
nitrogen may be obtained from certain compounds of the 
element. ' Thus if heat is applied to the compound, ammo- 
nium nitrite, NH4NO2, the change represented in the follow- 
ing equation takes place : 

NH.NO, = 2H,0 + 2N. 
Physical Properties. Nitrogen is similar to oxygen 
and, hydrogen in that is a colorless, odorless and taste- 
less gas. I liter of it weighs 1.2501 grams; hence it is 



Nitrogen and the Rare Elements, Etc. 81 

about 14 times as heavy as hydrogen and slightly lighter 
than oxygen. It can be obtained in the form of a colorless 
liquid having a boiling point of -195°. At -214° it be- 
comes a solid. 

Chemical Properties, Nitrogen is characterized by 
its inertness. It has but slight activity even when heated to 
high temperatures. _ It is neither combustible nor a sup- 
porter of combustion. At ordinary temperatures it will not 
combine directly with any of the elements. x\t higher tem- 
peratures it combines with magnesium, lithium, titanium 
and a few other elements. The compounds formed are called 
nitrides, just as compounds of an element with oxygen are 
■called oxides. When mixed with oxygen and subjected to 
the action of electric sparks combination slowly takes place 
with the formation of oxides of nitrogen. A mixture of 
nitrogen and hydrogen when treated similarly forms am- 
monia, a gaseous compound of nitrogen and hydrogen. 

Since we are constantly inhaling nitrogen, it is evident 
that it is not poisonous. Nevertheless life would be im- 
possible in an atmosphere of pure nitrogen on account of 
the exclusion of the necessary oxygen. 

Argon, Necn, IJelium, Krypton, Xenon. These are all rare 
•elements, recently discovered in the atmosphere, in which they 
occur in very small quantities. Argon, discovered in 1894, was- 
the first one obtained. Lord Rayleigh, an English scientist, 
while engaged in determining the exact weights of various 
gases, observed that nitrogen obtained from the air is slightly 
heavier than the pure nitrogen obtained from its compounds. 
After repeating his experiments many times, always with the 
same results, Rayleigh finally concluded that the nitrogen ob- 
tained from the air is not pure, but is mixed with a small 
amount of some unknown gas, the density of which is greater 
than that of nitrogen. Acting on this assumption, Rayleigh, 
together with the English chemist Ramsay, attempted to sep- 
arate the nitrogen from the unknown gas. Knowing that nitro- 
^gen would combine with magnesium, they passed the nitrogen 



82 Inorganic Chemistry. 

obtained from the air and freed from all known substances^ 
through tubes containing magnesium heated to the necessary 
temperature. After repeating this operation, they finally suc- 
ceeded in obtaining from the atmospheric nitrogen a small 
volume of gas which would not combine with the magnesium, 
and hence was not nitrogen. This proved to be a new element 
to which they gave the name argon. As predicted, this new 
element was found to be heavier than nitrogen, its density 
as compared with hydrogen as a standard being approximately 
20, that of nitrogen being only 14. About 1% of the atmos- 
pheric nitrogen proved to be argon. The new element is char- 
acterized by having no affinity for other elements. Even under 
the most favorable conditions, it can not be made to com- 
bine with any other elements. On this account it was given 
the name argon, signifying "lazy" or "idle." Like nitrogen,, 
it is colorless, odorless, and tasteless. It has been liquefied and 
solidified. Its boiling-point is — 187°. 

The remaining elements of this group have been obtained 
from liquid air. Under the subject of distillation attention 
was called to the fact that by repeated distillation a mixture 
of liquid^ of different boiling points can be separated from 
each other. In this way, by liquefying air and then allowing 
it to boil, those constituents having the lowest boiling, .points,, 
and hence the most difificult to liquefy, tend to escape first, 
followed by the others in the order of their boiling points. In 
this way Ramsay succeeded in obtaining from liquid air not 
only the known constituents, including argon and helium, but 
also the new elements neon, krypton and xenon. These new ele- 
ments all proved to be similar to argon in that they are with- 
out chemical activity, apparently forming no compounds what- 
ever. The amounts present in the air are very small. 

Helium was first found in the gases expelled from cer- 
tain minerals by heating. Through the agency of the spectro- 
scope it had been known to exist in the sun long before its 
presence on the earth had been recognized — hence the name 
helium, signifying the sun. It is the only gas which has not 
yet been liquefied, so that its boiling point must be below that of 
hydrogen. The names, neon, krypton, xenon, signify respec- 
tively, "new," "hidden," "stranger." 



CHAPTER VIII. 
THE ATMOSPHERE. 

Many references have already been made to the compo- 
sition and properties of the atmosphere. These statements 
must now he collected and discussed somewhat more in de- 
tail. 

The term atmosphere is applied to the gaseous envelope 
surrounding the earth. The term air is generally applied to 
a limited portion of this gas although often used synono- 
mously with the term atmosphere. 

Air Formerly Regarded as an Element. Like w^ater 
air was at first regarded as elementary in character.. Near 
the close of the eighteenth century, Priestley, Lavoisier and 
Scheele showed by their experiments that it is a mixture of 
at least two gases, those which we now call oxygen and 
nitrogen. By burning substances in an enclosed volume of 
air and noting the contraction in volume due to the removal 
of the oxygen, they were able to determine with some ac- 
curacy the relative volumes of these gases present in the air. 

The Constituents of the Atmosphere. 

I. Constituents Essential to Life. In addition to oxy- 
gen and nitrogen at least two other substances, namely, car- 
bon dioxide and water vapor must be present in the atmos- 
phere in order that life may exist. The former of these is a 
gaseous compound of carbon and oxygen having the formula 
CO2. Its properties will be discussed in detail in the chap- 
ter on the compounds of carbon. Its presence in the air 
may be shown by causing the air to bubble through a so- 
lution of calcium hydroxide, Ca(0H)2, commonly called 

(83) 



84 Inorganic Chemistry. 

lime water. The carbon dioxide combines with the calcium 
hydroxide in accordance with the following equation : 

Ca ( OH ) 2 + CO, = CaCOg + H^O 

The calcium carbonate, CaCOg, is insoluble in water 
and separates in the form of a white powder which causes 
the solution to appear milky. The presence of water vapor 
is readily shown by its condensation on cold objects as well 
as by exposing a bit of dry calcium chloride to the air, 
sufficient moisture soon being absorbed to dissolve the chlor- 
ide. 

Function of Each of the Essential Constituents. The 

oxygen directly supports life through respiration. The nitro- 
gen, on account of its inactivity, serves to dilute the oxygen, 
and while contrary to the older views, it is possible that life 
might continue to exist in the absence of the atmospheric 
nitrogen, yet the conditions of life would have to be entirely 
changed. Moreover, nitrogen is an essential constituent of 
all animal and plant life. It was formerly supposed that neither 
animals nor plants can assimilate the free nitrogen, but it has 
been shown recently that at least one order of plants, the 
leguminosiae, to which belong the beans and peas, has the 
power of directly assimilating the free nitrogen from the at- 
mosphere. 

The carbon dioxide in the air serves as a plant food. 
The plant absorbs it and decomposes it, utilizing the carbon 
in building up its tissues and liberating the oxygen. Water 
vapor is essential to both plants and animals, neither of which 
can exist for a long period in a perfectly dr}^ atmosphere. 

2. Constituents not Essential to Life. In addition to 
the oxygen, nitrogen, carbon dioxide and water vapor, the 
air contains small amounts of various other gases, the pres- 
ence of which so far as known is not essential to life. This 
list includes the rare elements argon, helium, neon, kryp- 
ton and xenon; also hydrogen, ammonia, hydrogen dioxide 
at^d probably ozone. Certain minute forms of life (germs) 



The Atmosphere. 85 

are also present, the putrefaction of organic matter being 
due to their presence. 

The Quantitative Analysis of Air. A number of dif- 
ferent methods have been devised for the determination of 
the amounts of the constituents of the atmosphere. The 
most important of which are the following. 
I. Determination of Oxygen. 

a. The oxygen is withdrawn from a measured vol- 
ume of air inclosed in a tube, by means of phosphorus. 
To make the determination, a graduated tube is filled 
with water and inverted in a vessel of water. Air is 
now introduced into the tube until it is partially filled with 
the gas. The volume of the inclosed air is carefully noted 
and reduced to standard conditions. A small piece of phos- 
phorus is attached to a wire and brought within the tube 
as shown in Fig. 26. After a few hours the oxygen in the 
enclosed air will have combined with the phosphorus, 
the water rising to take its place. The volume is 
again noted and reduced to standard conditions. The 
. contraction in the volume of the air is 
y^ equal to the volume of oxygen absorbed, 
b. The oxygen may also be removed 
from a measured volume of air by pass- 
ing it through a tube containing copper 
heated to a high temperature. The oxy- 
gen in the air combines with the copper 
to form copper oxide, CuO. Hence the increase in the 
weight of the copper equals the weight of the oxygen in 
the volume of air taken. 

c. A more accurate method is the following. 

A eudiometer tube is filled with mercury and inverted 

in a vessel of the same liquid. A certain amount of air is 

then introduced into the tube and the volume accurately 

noted. There is then introduced more than sufficient hy- 



86 Inorganic Chemistry. ^ 

drogen to combine with the oxygen present in the enclosed 
air, and the volume again accurately noted. The mixture 
is then exploded by an electric spark, and the volume is 
again taken. By subtracting this volume from the total 
volume of the air and hydrogen there is obtained the 
contraction in volume due to the union of the oxygen 
and hydrogen. The volume occupied by the water formed 
by the union of the two gases is so small in comparison 
with the volume of the gases which combined to form it 
that it may be disregarded altogether. It is evident then 
that the contraction in volume due to the combination is 
equal to the volume occupied by the oxygen in the air in 
the tube plus twice this volume of hydrogen; in other 
words one-third of the total contraction is equal to the vol- 
ume occupied by the oxygen in the enclosed air. The fol- 
lowing example will make this clear. 

Volume of air in tube -------50 c.c. 

Volume after introducing hydrogen - - - - 80 c.c. 

Volume of hydrogen introducedc=80 — 50 - - - -30 c.c. 

Volume after combination of oxygen and hydrogen - 48.5 c.c. 

Contraction in volume due to combination=80 — 48.5 31.5 c.c. 

Volume of oxygen in 50 c.c. of air=i: of 31.5 - - 10.5 c.c. 

The results obtained by these methods show that lOO 
volumes of air contain approximately 2i volumes of oxygen. 
2. Determination of Nitrogen. 

By passing the gas left after the removal of the 
oxygen over heated magnesium the nitrogen is with- 
drawn, argon and the other rare elements being left. It 
may thus be shown that of the 79 volumes of gas left after 
the removal of the oxygen from 100 volumes of air, approx- 
imately 78 are nitrogen and 0.93, argon, the other elements 
being present in such small quantities that they may be neg- 
lected. 



The Atmosphere. 87 

3. Deter uiination of Carbon Dioxide. 

The amount of carbon dioxide present in the air may 
be determined by absorbing it with calcium hydroxide or 
some similar compound, the increase in weight being equal 
to the carbon dioxide absorbed. The amount present in 
the open normal air is from 3 to 4 parts in 10,000. 

4. Dctennination of Water Vapor. 

The water vapor may be determined by absorbing it 
w^ith calcium chloride or some other good drying agent. 
The amount of it present varies not only with the locality 
but also varies greatly from day to day in the same locality, 
because of changes in temperature and the winds. 

Air a Mechanical Mixture. The results of a great 
many such analyses have shown that the composition of the 
air, apart from water vapor, is very nearly constant. Even 
when collected from widely differing localities, and at all 
seasons of the >ear, it is almost exactly the same in compo- 
sition, and would therefore appear to be a definite chemical 
comipound. If it is a mechanical mixture, its composition 
should vary from place to place, and from season to season, 
for there are many processes going on in nature which would 
tend to change the proportion of its constituents. 

That air is a mixture and not a compound can be shown 
in a number of ways among which are the following: 

1. Water dissolves oxygen and nitrogen from the air 
in an entirely different proportion from that in which they 
are present in the air. This would not be the case if the two 
were combined. 

2. A chemical compound in the form of a liquid has 
a definite boiling point. Water for example boils at 100°. 
Moreover the steam which is thus formed has the same 
composition as the water. The boiling point of liquid air 
on the other hand gradually rises as the liquid boils, the 



88 Inorganic Chemistry. 

nitrogen escaping first followed by the oxygen. If the two 
were combined, they would pass off together in the propor- 
tion in which they are found in the air. 

The surprising constancy in the composition of the air 
is explained by several facts. The winds keep the air in 
constant motion, and so prevent local changes. The volume 
of the air is so vast that any changes which do occur are 
very small compared with the total amount of the air, and 
cannot be noticed. The changes occasioned by living or- 
ganisms which are discussed in the next paragraph tend to 
balance each other to quite an extent. 

Changes in the Atmosphere due to Respiration and 
Combustion. In the process of respiration, the oxygen 
in the inhaled air is absorbed by the blood and carried to 
all parts of the body, where it combines with the carbon of 
the worn out tissues. The products of oxidation are car- 
ried back to the lungs and exhaled in the form of carbon 
dioxide. All the ordinary forms of fuel contain large per- 
centages of carbon. On burning, this carbon combines^ with 
the oxygen in the air forming carbon dioxide. Evidently 
in these two processes large amounts of oxygen are with- 
drawn from the atmosphere and returned to it in the form of 
carbon dioxide. This in itself would cause a diminution in 
the percentage of free oxygen present. 

Plants however have the power, in the sunlight, of ab- 
sorbing this carbon dioxide, retaining the carbon and return- 
ing the oxygen once more to the. air, so that these two pro- 
cesses largely balance each other. 

While air collected in the open fields shows a practically 
uniform composition, the case is different when it is collected 
from a closed or poorly ventilated room, occupied by a 
number of persons. Such air shows a rapid rise in the 
amount of carbon dioxide present. While this gas is not 
to be regarded as poisonous unless present in large amounts. 



The Atmosphere. 89 

nevertheless air containing more than 15 parts in 10,000, 
is not fit for respiration. 

Along with the carbon dioxide a certain amount of or- 
ganic matter is exhaled, which imparts to the respired air 
its foul odor. It is probable that this organic matter rather 
than the carbon dioxide is responsible for the injurious ef- 
fects which follow the respiration of impure air. The ex- 
tent of such impurities present may be judged however from 
the amount of carbon dioxide present, since the two are ex- 
haled together. 

The Properties of the Atmosphere. Inasmuch as the 
atmosphere is composed principally of a mixture of oxygen 
and nitrogen, which elements have already been discussed, 
its properties may largely be inferred from those of the two 
gases. I liter of it weighs 1.292 grams. It is thus 14.4 
times as heavy as hydrogen. At the sea level it exerts an 
average pressure suffici-ent to sustain a column of mercury 
760 mm. in height. This is taken as the standard pressure 
in determining the weights of gases as well as the boiling 
points of liquids. Water may be made to boil at any tem- 
perature between 0° and considerably above 100° by simply 
varying the pressure. It is only when the pressure upon it 
is equal to the normal pressure of the atmosphere at the sea 
level as indicated by a barometric reading of 760 mm. that 
it boils at 100°. 

Liquid Air. Attention has already been called to 
the fact that both oxygen and nitrogen can be obtained in 
the liquid state by strongly cooling the gases and applying 
great pressure to them. Air being a mixture of these two 
gases can naturally be liquefied by the same methods. 

Recently the methods for liquefying air have been much 
simplified in that artificial cooling is dispensed with, the 
temperature being reduced sufficiently by allowing a por- 
tion of the compressed air to expand. The expansion of a 



90 



Inorganic Chemistry. 




Fig. 27 



gas is always attended by the absorption of heat. In lique- 
fying air the apparatus is so constructed 
that the heat absorbed is withdrawn from 
air already under great pressure. This 
process is continued until the temperature 
is lowered to the point of liquefaction. ' 

It is not possible to preserve air in a 
liquid state by confining it, on account of 
the enormous pressure exerted by it in its 
tendency to pass into the gaseous state. It 
may be preserved for some hours or even 
days before it will completely evaporate, by 
simply placing it in an open vessel surrounded by a non- 
conducting material. The most efficient vessel for retain- 
ing it is the "Dewar bulb" shown in Fig. 27. The air is 
withdrawn from the space between the two walls thus mak- 
ing it non-conducting. One liter of liquid air placed in such 
a bulb m.ay be preserved nearly two weeks before complete 
evaporation takes place. 

When first prepared, liquid air is cloudy because of the 
presence of particles of solid carbon dioxide. These may 
be filtered off, leaving a liquid of slightly bluish color. It 
begins to boil at about — 190°, the nitrogen passing off first 
gradually followed by the oxygen, the last portions bemg 
nearly pure oxygen. In this way oxygen is now prepared 
commercially. 

Liquid air is also used as a cooling agent in the study 
of the properties of matter at low temperatures. It has thus 
been found that elements at extremely low temperatures 
largely lose their chemical activity. 

The extremely low temperature of liquid air may be 
inferred from the fact that mercury when subjected to its 
action is frozen to a mass so hard that it may be used for 
driving nails. 



CHAPTER IX. 




SOLUTIONS. 

It has been seen that when a substance disappears in a 
liquid in such a way as to thoroughly mix with it and be 
lost to sight as an individual body, the resulting product is 
called a solution. We shall frequently meet with two kinds 
of solutions and a brief study of them will be of advantage 
at this point. 

I. Solution of Gases in Liquids. It has already 
been stated that oxygen, hydrogen and nitrogen are slightly 
soluble in water; accurate study has led to the conclusion 

that all gases are soluble to 
some extent not only in 
water, but in many other 
liquids. The amount of a 
gas which will dissolve in a 
liquid depends upon a num- 

, ber of conditions, and these 

=-:i:J^ can best be understood by 
--=-^\ supposing the gas to be con- 
fined over the liquid in such 
a way that it cannot escape or be displaced by another gas. 
Fig. 28 represents a gas confined in this way in a bell jar, 
which is inverted and placed over a liquid. The following 
conditions will determine how much of the gas will dissolve 
in the liquid. 

(a) Nature of the Gas, Other conditions being equal 
some gases are much more soluble than others, so that each 
gas has its own peculiar solubility just as it has its own 
especial taste or odor. The solubility of gases varies be- 
tween wide limits as will be seen from the following table, 

(91) 



Fig. 28 



92 Inorganic Chemistry. ^ 

but as a rule a given volume of a liquid will not dissolve 
more than two or three times its own volume of a gas. 

SOLUBILITY OF GASES IN WATER. 

i,ooo c.c. of water at 760 mm. pressure and at 0° will 
dissolve 

Ammonia 1148. Hters 

Hydrochloric acid 503 . liters 

Sulphur dioxide 79-79 liters 

Carbon dioxide 1.8 liters 

Oxygen 41 • 14 c.c. 

Hydrogen 21.15 c.c. 

Nitrogen 20. 03 c.c. 

In the case of very soluble gases, such as the first three 
in the table, it is probable that feeble chemical combination 
between the liquid and the gas takes place. 

(b) Nature of the Liquid. The character of the Jiquicf 
has much influence upon the solubility of a gas. Water^ 
alcohol, ether, each has its own peculiar solvent power. From 
the solubility of a gas in water, no prediction can be made 
as to its solubility in other liquids. 

(c) Influence of Pressure. It is found that the amount 
of gas which dissolves in a given case is proportional to the 
pressure exerted upon the gas. If the pressure is doubled, 
the amount of gas going into solution is doubled ; if the 
pressure is diminished to one-half of its original value, half 
of the dissolved gas will escape. Under high pressure, large 
quantities of gas can be dissolved in a liquid, and when the 
pressure is removed the gas escapes, causing the liquid to 
foam or eflfervesce. It is the sudden release of the pressure 
upon liquids containing carbon dioxide which causes the 
efifervescence of soda water and similar substances. 



Soliifiojis. 93 

(d) Influence of Temperature. In general, the lower 
the temperature of the liquid, the larger the amount of gas 
which it is able to dissolve. looo volumes of water at o^ 
will dissolve 41.14 volumes of oxygen; as the temperature 
rises, the amount of oxygen remaining in solution becomes 
less and less until at the boiling point all the oxygen is ex- 
pelled. 

II. Solutions of Solids in Liquids. Physical and 
and Chemical Solutions. The distinction between chemical 
and physical solutions must be carefully noted. A\'e speak 
of dissolving zinc in sulphuric acid and likewise of dissolv- 
ing salt in water, but these two acts are quite different. On 
evaporating the solution of salt, the salt is recovered en- 
tirely unchanged: by evaporating the solution of zinc in 
sulphuric acid it is impossible to regain the zinc. In its 
place is obtained a soft, white substance, bitter in taste, 
easily soluble in water — in short with properties entirely 
different from those of the original zinc. 

The solution of the zinc was preceded by a chemical 
change which converted the zinc into another substance, 
zinc sulphate, which then dissolved in the water. Solutions 
of this kind are called chemical solutions, because accom- 
panied by chemical changes in the substance dissolved. So- 
lutions like that of salt in water, in which the dissolved sub- 
stance can be recovered unchanged by evaporating the sol- 
vent, are called physical solutions, and it is with them that 
we are noAv concerned. 

Circumstances Affecting the Solubility of a Solid. 

(a) Xafure of the Solid. Other conditions being the 
same, solids vary greatly in their solubility in liquids. Some 
are very soluble, others moderately so, while still others are 
said to be insoluble. Thus 100 parts of water at 18^ will 
dissolve : 



94 



Inorganic Chemistry. 



Calcium chloride 71 . parts 

Sodium chloride . . . . 35-9 parts 

Potassium nitrate 29 . i parts 

Copper sulphate 21 .4 parts 

Calcium sulphate 0.207 parts 

None are absolutely insoluble,- but the amount dissolved 
may be so small as to be of no significance for most pur- 
poses. Thus barium sulphate, one of the most insoluble of 
common substance, dissolves in water to the extent of i 
part in 400,000. 

(b) Nature of Solvent. Liquids vary much in their 
ability to dissolve solids. Some are said to be good solvents, 
since they dissolve a great variety of substances, and consid- 
erable quantities of them. Others have small solvent power, 
dissolving few substances, and those to but a slight extent. 
Broadly speaking, water is the most general solvent, and al- 
cohol is perhaps second in solvent ability. 

(c) Temperature. The amount of solid which a given 
liquid can dissolve varies with the temperature. Usually it 
increases rapidly as the temperature rises, so that the boil- 
ing liquid dissolves several times the amount which the 
cold liquid will dissolve. In some instances, as in the case 
of common salt dissolved in water, the temperature has 
little influence upon the solubility, and a few solids are more 
soluble in cold water than in hot. The following examples 
will serve as illustrations : 



TABLE OF SOLUBILITY AT C 


) AND AT TOO . 


parts of water will dissolve 




at 0° 


at 100° 


Calcium chloride .... 49.6 


... 155. parts 


Sodium chloride .... 35.7 


39.8 parts 


Potassium nitrate ... 13.3 


. . . 247. parts 


Copper sulphate .... 15.5 


. . . 73.5 parts 


'Calcium sulphate .... 0.205 


0.217 parts 



Solutions. ■ 95 

Saturated Solutions. A liquid will not in general 
dissolve an unlimited amount of a solid. It becomes satu- 
rated, and when this point is reached no more of the solid 
will dissolve at that temperature. It is possible therefore 
to determine accurately the solubility of a substance for any 
given temperature. The liquid will however dissolve any 
quantity of the solid which is less than the amount required 
for saturation, and in this respect a solution differs very 
strikingly from a chemical compound in which, as we have 
seen, the constituents combine in certain definite proportions 
and in no others. 

Crystallization. When a hot, saturated solution is 
allowed to cool, some of the dissolved solid must in general 
separate from the solution ; for as we have seen, solids are 
usually more soluble in hot liquids than in cold. Usually 
the substance separates in the form of solid particles of 
definite geometric form, called crystals, and the process just 
described is called crystallization. Sometimes it separates as 
a powder which may consist of minute crystals, or may be 
without crystalline form. In the latter case it is said to be 
amorphous. 

General Physical Properties, (a) Distribution of the 
Solid in the Liquid. A solid, when dissolved, tends to 
distribute itself uniformly through the liquid so that every 
part of the solution has the same strength or concentration. 
This process goes on very slowly unless hastened by stirring 
or shaking the solution. Thus if a few crystals of a highly 
colored substance such as potassium permanganate are 
placed in the bottom of a tall vessel full of water, it will 
be seen that it takes a long time for the solution to become 
uniformly colored. 

(b) Boiling Point of Solutions. The boiling point 
of a liquid is raised by the presence of a substance dissolved 
in it. In general the amount of elevation of the boiling 



96 Inorganic Chemistry. 

point of a liquid, caused by a given substance, is proportional 
to the amount of the soUd dissolved. The extent of this 
elevation may be considerable when the solid is very soluble. 
A saturated solution of calcium chloride for example boils 
at 179.5°. 

(c) Freezing Point of Solutions. A solution freezes 
at a lower temperature than the pure solvent. The lower- 
ing of the freezing point obeys the same law which holds 
for the raising of the boiling point — the amount of lower- 
ing is proportional to the amount of dissolved substance, 
that is to the concentration of the solution. The fact that 
sea water is usually free from ice when fresh water near by 
is frozen over is in part due to the large amount of salt 
which the sea water contains. 

Electrolytic Dissociation. A large number of facts 
which can not be taken up in an elementary course have led 
chemists to believe that many substances when in solution 
imdergo an important change. A portion of their molecules 
fall apart or dissociate into two or more parts which have 
been named ions. Sodium nitrate,- NaNOg, for example dis- 
sociates into the ions Na and NO3. These ions differ 
widely in properties from the elements or groups of elements, 
in that they are heavily charged with electricity. Thus in 
the dissociation of NaNOg, the Na ion carries a positive 
charge of electricity while the NO3 ion carries a negative 
charge. Moreover the positive charge carried by the Na 
ion is exactly equal to the negative charge carried by the 
NO3 ion so that the solution as a whole is electrically neu- 
tral. Certain molecules may dissociate into more than two 
ions. Thus sodium sulphate, Na2S04, dissociates into two 
Na ions and one SO4 ion. In this case the positive charges 
carried by the two Na ions are exactly equal to the negative 
charges carried by the SO4 ion. In the dissociation of com- 
pounds of the metals it is always the metal ion which car- 
ries the positive charge. 



Solutions. 



97 



The dissociation of compounds is sometimes represented 
in the form of equations like the following which represents 
the dissociation of sodium nitrate: 



/ NaN03 = Na + NO3 

The marks + and — indicate the character of the 
■charges carried by the ions. 

These ions are free to move about in a solution inde- 
pendently of each other and for this reason were given the 
name "ion" which signifies "a wanderer.'' 

Electrolytes. Solutions differ much in their ability 
to conduct the electrical current. Some, such as a solution 
of sugar in water, do not conduct the current appreciably, 
Avhile others, as for example a solution of salt in water, 
•conduct it vrith great ease. A substance whose solution will 
conduct the current is called an electrolyte, and a close study 
of these electrolytes has shown that they are all substances 
Avhich dissociate into ions in solution. This fact has sug- 
•gested that it is the ions which carry the current through a 
■solution, and that consequently a solution containing no 
ions cannot conduct the current. 

Electrolysis. When an electrical current is passed 
through the solution of an electrolyte, some chemical change 
;always takes place in the electrolyte and often in the solvent. 
This change is called, electrolysis. The dissociation theory 
•oflFers a simple explanation of the facts of electrolysis which 
ivill be easily understood by reference to the following dia- 
-gfram : 




Fig, 29 



98 



Inorganic Chemistry. 



When two electrodes, A and B in Fig. 29, are connected 
with the poles of a battery and placed in a solution of an 
electrolyte, the positive ions, called cations, are attracted 
to the negative electrode (cathode) and the negative ions, or 
anions, are attracted to the positive electrode, (anode). On 
reaching the electrodes the ions are discharged and have then, 
the properties of ordinary chemical substances. Thus if the 
electrolyte is sodium chloride dissolved in water, as repre- 
sented in Fig. 29, the sodium ions on being discharged 
immediat-ely attack the water, just as a bit of the metal so- 
dium would do, and hydrogen is set free according to the 
equation Na + H.O = NaOH + H 

The chlorine ion, on being discharged, may either be 
given off as chlorine gas or may act upon the water as rep- 
resented in the equation 

2 CI + H,0 = 2 HCl + O 
In a similar way, sodium sulphate, Na2S04, gives the ions 
2Na and SO4. On being discharged, the sodium atoms de- 
compose water about the cathode as indicated above while the 
SO4 ions when discharged at the anode decompose tYit water 
as represented in the equation 

S04 + H20 = H2S04+0 
That NaOH is formed at the cathode and H2SO4 at the 
anode may be demonstrated in the following way: NaOH 
has the power to turn red litmus blue,, 
while H2SO4 turns blue litmus red. A 
U-tube, such as is represented in Fig. 
30, is partially filled with a solution of 
sodium sulphate, and the liquid in one 
arm is colored with red litmus, that in 
the other with blue litmus An elec- 
trode placed in the red solution is made 
to serve as cathode, while one in the 
blue solution is made the anode. On al- 
lowing the current to pass, the blue so- 
turns red, indicating the formation of 
Fig. 30. H2SO4 at the anode, while the red so- 

tion turns blue, indicating the presence of NaOH at the cathode. 




Solutions. 99 

Electrolysis of Water. The reason for the addition 
of sulphuric acid to water in the preparation of oxygen and 
hydrogen by electrolysis will now be clear. Water itself is 
not an electrolyte, that is it does not form enough ions to 
carry a current. Sulphuric acid is an electrolyte, and in so- 
lution in water, dissociates into the ions 2 H and SO4. 

In the process of electrolysis of the solution, the hydro- 
gen ions travel to the cathode and on being discharged es- 
cape as hydrogen gas. The SO^ ions when discharged at 
the anode, act upon water, setting free oxygen and once 
more forming sulphuric acid. 

SO, + H,0 = H,SO, + O 

The sulphuric acid can again dissociate, and the process 
repeat itself as long as any water is left. Hence the hydro- 
gen and oxygen set free in the electrolysis of water really 
come directly from the acid, but indirectly from the water. 

Although the theory of electrolytic dissociation may at 
first sight appear to be fanciful, it must be remembered that 
it is supported by a great many experimental facts, and it 
has very successfully stood the many tests which have been 
applied to it. The theory has been very useful in suggesting 
an explanation of many facts of chemistry which were for- 
merly difficult to explain. It has also been of much value 
in leading to the discovery of new facts. In both respects 
therefore it has well served the purpose of a theory. We 
shall meet with many applications of the theory in subse- 
quent pages. 



Lof 



[ CHAPTER X. : . ■ 

ACIDS, BASES, SALTS, VALENCE. 

Acids. It is impossible to advance very far in the study 
of chemistry without meeting with substances called acids, 
and repeated mention has already been made of represen- 
tatives of this class. The acids most frequently used in 
the laboratory are hydrochloric, nitric and sulphuric acids. 
These will be described in detail in subsequent chapters, 
but a few words concerning them are necessary at this place. 

Hydrochloric acid is a gas composed of hydrogen and 
chlorine and has the lormula HCl. It is very soluble in 
water, and it is this solution which is usually termed hydro- 
chloric acid. Nitric acid is a substance composed of hy- 
drogen, nitrogen and oxygen having the formula HNO3. 
It is a heavy liquid, ind as sold commercially contains about 
32% of water. Sulphuric acid, whose composition is' repre- 
sented by the formula H0SO4, is an oily liquid nearly twice 
as heavy as water, and is commonly called oil of vitriol. 

Properties Common to Acids. Acids have a num- 
ber of properties in common, (i) They all contain hydro- 
gen. (2) In dilute solution they have a sour taste recalHng 
that of vinegar. (3) They have the property of changing 
the color of certain substances. Thus a solution of the blue 
coloring matter called litmus is turned red by even a trace 
of an acid ; methyl orange is, in a similar way, changed 
from red to yellow. (4) The most important property of 
acids however, is their power to act chemically upon sub- 
stances called bases. In this action the characteristic prop- 
erties of both acid and base disappear, and substances with 
properties quite different from either of them are formed. 

(100) 



Acids, Bases, Salts, Valence. loi 

Bases. Among the members of this class commonly 
used in the laboratory are sodium hydroxide, NaOH, po- 
tassium hydroxide, KOH, and calcium hydroxide, 0( OH) 2. 
These are white solids soluble in water, the calcium hydrox- 
ide being only sparingly so. Not all bases however are sol- 
uble in water. The strongest bases are sometimes called 
alkalies. 

Properties Common to Bases, (i) Bases consist 
of a metallic element such as sodium, calcium, copper, iron, 
in combination with hydrogen and oxygen. (2) When sol- 
uble in water, they have a disagreeable, brackish taste. (3) 
They reverse the color change produced by acids, restoring 
the blue color to red litmus and the red color to methyl 
orange. (4) They have the power to act upon acids and 
in this reaction the peculiar properties of both acids and bases 
are lost. 

Salts. When an acid acts upon a base, the hydro- 
gen of the acid combines with the oxygen and hydrogen of 
the base to form water, while the remainder of the acid con- 
stituent combines with the metal of the base to form a sub- 
stance called a salt. Thus sodium hydroxide acts with hy- 
drochloric acid to form the salt sodium chloride, NaCl, and 
water as follows : 

NaOH + HCl = NaCl + H^O 

Similarly calcium hydroxide acts with nitric acid to 
form the salt calcium nitrate, Ca(N03)2 and water, the 
equation being, 

Ca(OH)2 + 2 HNO3 = Ca(N03)2 -f 2 H,0 

This action of a base and an acid with each other is 
called neutralization. The products of neutralization are al- 
ways a salt and water. 

Properties Common to Salts, (i) A salt consists of 
a metal combined with the constituents of an acid other than 
its hydrogen. It may thus be regarded as derived from an 



I02 Inorganic Chemistry. 

acid by replacing the hydrogen of the acid with a metal. (2) 
They do not in general change the color of either red or blue 
litmus and for this reason are said to be neutral. (3) They 
have usually a salty or bitter taste quite different from the 
characteristic taste of acids and bases. 

It is usually true that a base contains an equal number 
of atoms of oxygen and hydrogen and since these two elements 
nearly always act together in any chemical action which the 
base undergoes, the symbols of these elements are grouped 
together in writing the formula of the base. Thus the formula 
of calcium' hydroxide is written Ca(OH)2 rather than Ca02H2. 
Any group of elements which act together in this way is called 
a radical. Some of these radicals have been given names, the 
name signifying the elements present in the radical. Thus the 
group OH is called the hydroxyl radical. 

Neutralization and the Dissociation Theory. The 

theory of electrolytic dissociation affords an interesting ex- 
planation of the facts of neutralization. Acids, bases and 
salts are all good electrolytes ; that is their solutions are all 
good conductors of electricity. They are therefore sub- 
stances which part, or dissociate, into ions when in solu- 
tion in water. When acids dissociate the positive ion is al- 
ways hydrogen, while the composition of the negative ion 
depends on the composition of the acid. Nitric, hydrochloric 
and sulphuric acids dissociate as indicated in the formulas; 

+ - + - + 

H, NO3 H, CI 2H, SO4 

Bases dissociate into a positive metallic ion, and one or 

more negative hydroxyl ions as indicated in the formulas ; 

+ - ++ - 

Na, OH Ca, 2(0H) 

In accordance with this view an acid may be defined 

as a substance which gives hydrogen ions in solution, while 

a base is a substance which gives hydroxyl ions in solution. 

Water is not appreciably dissociated, and whenever hydro- 



Acids, Bases, Salts, J^aleuce. 103 

gen and hydroxyl ions come together they unite to form 
iindissociated water. Hence when solutions of an acid and 
"base are mixed, the positive hydrogen ion of the acid and 
the negative hydroxyl ion of the base combine to form 
^lectricahy neutral water, and the other ions remain un- 
changed. On driving oil the Avater of the solution by 
evaporation, thev gradually unite to form the neutral salt. 
The cause of the reaction is thus seen to be the formation 
of undissociated water from hydrogen and hydroxyl ions. 

Metals. The elements fall naturally into two 
•classes : ( i ) those which in combination with hydrogen and 
oxygen, form acids, and (2) those which, in combination 
with the same elements, form bases. The former are called 
non-metals ; the latter, metals. ^lany of the metals are 
familiar substances, such as iron, copper, tin. They will be 
taken up for study after the non-metals have been con- 
sidered. At present it need only be noted that they are a 
class of elements capable of forming bases when combined 
with hydrogen and oxygen, and of giving a positive ion 
wdien in solution. 

Neutralization a Definite Act. The act of neutrali- 
zation has been represented in a preceding paragraph by 
means of equations, and we have seen that equations should 
only be used to express a quantitative action between sub- 
stances. The question therefore arises : Is neutralization a 
quantitative action between fixed amounts of acid and base? 
]\Iany careful experiments have shown that it is. A definite 
amount of a base always requires exactly the same amount 
of a given acid for its neutralization. This can be demon- 
strated by weighing the substances which act upon each 
other ; but it is more convenient to "measure the volume of 
solutions of acid and base which are required for neutrali- 
zation, having previously determined the strength of each 
solution. The experiment can be carried out as follows : 



104 



Inorganic Chemistry. 



CM 



02^ 



Fig. 31 



Prepare a solution of sodium hydroxide by diluting 20 
c.c, of a 10% solution (the ordinary laboratory reagent) to 
100 c.c; also a solution of sulphuric acid by 
adding i c.c. of the strong acid to 100 c.c. of 
water. Fill two clean burettes with the twO' 
solutions, supporting them as shown in Fig. 31. 
By means of the stopcocks, draw off the solu- 
tion in each burette until the level of the liquid 
is just at the zero mark. Now let exactly 15. 
c.c. of the acid flow into a small beaker, add a 
few drops of neutral solution of litmus, and 
then run in the alkali, drop by drop, stirring^ 
constantly with a glass rod. The instant the so- 
lution turns blue, stop adding the alkali, and 
note the amount which has been added. Re- 
peat the experiment, using different amounts 
of the acid, say 10 c.c. and 20 c.c. Calcu- 
late in each case the amount of alkali re- 
required to neutralize i c.c. of the acid. 
It will be found that the same amount 
will be required in each case, showing that 
the act of neutralization is a definite one. If the strengths of 
the two solutions were accurately known, the amouaits by 
weight of the acid and base could be readily calculated. 

Replacing Power of the Metals. The neutralization 
of some of the common acids by two of the most frequently- 
used bases may be represented by the following equations: 
NaOH + HCl = NaCl + H^O 
NaOH + HNO3 = NaNOe + H^O 
2NaOH -f H2SO4 — Na^SO^ + 2H2O 
Ca(OH)2 + 2HCI = CaCl^ + 2H2O 
Ca(OH)2 + 2 HNO3 = Ga(N03)2 + 2 H^O. 
Ca(OH)2 + H2SO4 = CaSO^ -f 2 H^O 
A study of these reactions will show a singular fact. It 
appears that one atom of sodium can replace but a single 
atom of hydrogen in an acid, while an atom of calcium can 
replace two. A study of the metallic elements shows that 
they differ much in the number of hydrogen atoms which 



Acids, Bases, Salts, Valence. 105 

they are able to displace in an acid. Some are like sodium, 
and can take the place of but a single hydrogen atom ; some, 
like calcium, can replace two ; still others can replace three. 
This peculiar property of the metals is called their replacing 
power. We have no knowledge as to why the atoms differ 
in this respect, and there is no way to tell how many hydro- 
gen atoms a given metal will replace save by experiment. 

Replacing Power of Metals and the Formulas of 
Salts, When the replacing power of a metal and the for- 
mula of an acid are known it is easy to predict the prob- 
able formula of the salt which the metal will form with the 
acid. Thus sodium can replace but a single hydrogen atom ; 
consequently with the acids HCl, H2SO4, H3PO4, it gives 
the salts NaCl, NagSO^, NagPO^. Calcium can replace two 
hydrogen atoms, and the formulas of the calcium salts of the 
same acids will be CaClg, CaS04, Ca3(P04)2. 

In writing the equation expressing the neutralization 
of an acid by a base, it is therefore necessary to take the two 
in such proportions that the replacing power of the metal 
atoms of the base will equal the hydrogen atoms of the acid. 
The correct formula of the salt will then be easily obtained. 

Valence. Elements differ not only in the number of 
hydrogen atoms which they can displace from an acid, but 
also in the number of atoms of other- elements which they 
are able to hold in combination. Thus we have the com- 
pounds represented by the formulas : 

HCl, HoO, H3N, H4C. 

H^O, CaO, AI2O3, CO2, N2O5, SO3. 

It will be noticed that elements differ much in respect 
to the number of hydrogen atoms which they can hold in 
combination, and the same difference appears in regard to 
the oxygen atoms, as is seen in the formulas of the oxides. 

That property of an element which determines the num- 
ber of atoms of another element which it can hold in com- 



io6 Inorganic Chemistry. ^ 

• 
bination, is called its valence. Valence is strictly a numerical 
relation between the elements, and does not tell anything 
about the strength of the attraction between them. Thus 
oxygen has a far stronger attraction for hydrogen than does 
nitrogen, but it can only combine with two atoms of hydro- 
gen whereas nitrogen can combine with three. 

Measure of Valence. In expressing the valence of 
an element we must select some standard of value just as 
in the measure of any other numerical quantity. If we se- 
lect the oxygen atom as a standard of valence, a glance at 
the formulas of the oxides given above will show that hy- 
drogen will have a valence of J, calcium of i, aluminium of 
ij. The atom of hydrogen never combines with more than 
one atom of another element, and therefore serves as a better 
standard. It is called a univalent element. That property of 
the metals which we have called their replacing power is 
thus seen to be their valence as indicated by the number of 
hydrogen atoms which an atom of the metal can replace in 
an acid. 

Elements like chlorine which combine with hydrogen 
to form compounds consisting of a single atom of each, 
such as HGl, are likewise said to be univalent. Elements 
such as oxygen, which combine with two hydrogen atorns, 
forming such compounds as H2O, are said to be divalent; 
and in like mariner there are elements which are trivalent, 
tetravalent, pentavalent. None have a valence of more 
than eight. 

Many elements, especially among the metals, do not 
readily form compounds with hydrogen, and their valence 
is not easy to determine directly. They will, however, com- 
bine with other univalent elements, and their valence can be 
determined from these compounds. 

It will not do to assume that because the molecule of a 
compound is made up of a single atom of each of two ele- 



Acids, Bases, Salts, Valence. 107 

ments, that each of these is univalent. Thus we have the com- 
pound CaO ; but the formulas of the substances CaCl2 and 
HgO show that both calcium and oxygen are divalent. 

Variable Valence. Many elements are able to exert 
different valences depending on circumstances. Thus we 
liave the compounds CuoO and CuO ; CO and CO^ ; FeClg 
and FeClo. It is not always possible therefore to assign a 
:fixed valence to an' element; but each element lends to 
exert some normal valence, and compounds in which it has 
a valence different from this are apt to be unstable and easily 
-changed into the compounds of normal valence. 

Structure of Compounds. It often happens that we 
cannot decide upon the valence of an element in a compound 
imtil we know how the atoms are arranged in the molecule. 
Thus in the compounds water and hydrogen dioxide, H2O 
^nd H2O2, the oxygen appears to have the valences 2 and i. 
If we know that the atoms are arranged as represented in 
the formulas H-O-H and H-O-O-H, it is evident that the 
'Oxygen has in each case a valence of two. 

In such a compound as sulphuric acid, H2SO4, we 
must know much more than the mere formula tells us be- 
fore we can draw conclusions as to the valence of the 
.•elements in it. If it is known that the atoms are arranged 
as represented in the formula 5q> S ^S, the valence of the 
oxygen being two and that of the hydrogen one, the valence 
of the sulphur will be six. ' 

It is quite possible to gain such information about com- 
pounds, and the formulas which express this knowledge are 
•called structural formulas. The structural formula for sul- 
phuric acid has been shown by many careful experiments 
to be ^Q> S ^5. The elements whose symbols are joined 
by lines are supposed to be directly combined with each 
other, and the number of lines leaving the symbol of an ele- 



io8 Inorganic Chemistry. ^ 

ment, indicates the valence of the element. In a similar way 
calcium Oxide may be represented by the formula Ca = O ; 
sodium hydroxide by the formula Na-O-H ; nitric acid by 
the formula H— 0-N ^ ^ 

It must constantly be remembered that such formulas 
have value only as they represent facts which have been 
established by experiment. It should be carefully noted also 
that valence itself is not a force or an attraction, and that 
two lines between two atoms, as in the formula Ca=0, does 
not mean that they are held together with double the at- 
traction which unites the elements in the compound H-Cl, 
but merely that each of the elements is divalent. 

Nomenclature of Acids, Bases and Salts. Acids, 
bases and salts are named in accordance with a few simple 
principles, the names being suggested by the characteristic 
elements in the compounds. The acids and salts are, for 
the purpose of naming, grouped in two classes ; binary, or 
those containing but two elements, and ternary, or those 
containing three. Some contain more than three,' but we 
shall meet very few of them. 

Nomenclature of Binary Acids. A binary acid is named 
by combining the names of the elements composing it, so 
as to form a euphonious word, the termination ''ic" being 
used. The word hydrogen is always abbreviated to ''hydro,'* 
which is used as a prefix. Thus we have HCl, hydrochloric 
acid ; H2S, hydrosulphuric acid. 

Nomenclature of Ternary Acids. In addition to the 
hydrogen present in all acids, nearly all of the ternary 
acids which we shall study also contain oxygen. In ad- 
dition to the hydrogen and oxygen, each ternary acid con- 
tains a third element which gives its name to the acid, the 
termination "ic" being used. Thus HNO3 is called nitric 
acid ; H2SO4 is called sulphuric acid ; HCIO3 is called 
chloric acid. 



Acids, Bases ^ Salts, Valence. 109 

There may be more than one acid, however, composed 
of the same elements, but differing in the amounts of 
oxygen present. If two such acids are known, the one 
containing the smaller amount of oxygen has the termina- 
tion "ous." Thus we have: 

H2SO4 sulphuric acid. HNO3 nitric acid. 

H2SO3. . .sulphurous acid. HNOo nitrous acid. 

In some cases more than two acids are known differ- 
ing only in the amount of oxygen present. In such cases 
the one containing less oxygen than the ''ous" acid, takes 
the name of the "ous" acid with the prefix *'hypo," while 
the one containing more oxygen than the "ic" acid, takes 
the name of the *'ic" acid with the prefix "per." Thus the 
names and formulas of the four acids containing hydro- 
gen, chlorine and oxygen are 

HCIO . .hypochlorous acid. HCIO3 chloric acid. 

HCIO2 chlorous acid. HiC104 . . . .perchloric acid. 

Nomenclature of Binary Salts. A binary salt is named 
from the elements composing it, the name of the non-metal 
receiving the termination "ide." Thus we have 

NaCl sodium -chloride. AggS silver sulphide. 

K I potassium iodide. CaFg calcium fluoride. 

Nomenclature of Ternary Salts. It has been stated that 
a salt may be regarded as derived from an acid by substi- 
tuting a metal for its hydrogen. Ternary salts receive their 
names from the acids from which they may be regarded 
as derived, by changing the termination "ic" of the acid 
to "ate;" or in case it is derived from an "ous" acid, the 
"'ous" is changed to "ite," the name of the metal being pre- 
fixed in either case. This method is illustrated in the follow- 
ing table, which gives the names and formulas of the ternary 
chlorine acids as well as the names and formulas of their 
corresponding sodium salts: 



no Inorganic Chemistry. ^ 

HCIO hypochlorous acid NaClO sodium hypochlorite 
HCIO2 chlorous acid NaClOo sodium chlorite 

HCIO3 chloric acid NaClOg sodium chlorate 

HCIO^ perchloric acid NaC104 sodium perchlorate 

Nomenclature of Bases. The naming of bases or hydrox- 
ides, presents few difficulties, the name being made up of 
the name of the metal and the word hydroxide, e. g. sodium 
hydroxide, (NaOH). Some metals are able to form more 
than one hydroxide, since they can exercise two different 
valences. The base with the smaller number of hydroxyl 
groups adds the suffix -ous to the name of the metal, while 
the one with the larger number of hydroxyl groups adds 
-ic. Thus we have cuprous hydroxide, Cu(OH), and cupric 
hydroxide, Cu(OH)2; ferrous hydroxide, Fe(OH)2, and 
ferric hydroxide, Fe(OH)3. 

Compounds of the metals, other than the hydroxides, are 
named in the same way, the sufflx -ous indicating a smaller 
valence than the suffix -ic. Thus we have cuprous chloride, 
CuCl, and cupric chloride, CuCls; cuprous oxide, Cuj^, and 
cupric oxide, CuO. 

Acid, Basic and Normal Salts. It has been seen that 
in neutralization the hydrogen of an acid is replaced by a 
metal. When an acid has more than one hydrogen atom 
in it however, it is possible with proper care to replace only 
a part of the hydrogen. Thus sodium may replace only 
one of the hydrogen atoms of sulphuric acid, forming a 
compound of the formula NaHS04. 

NaOH -f H2SO4 = NaHSO^ -|- H^O 

Such a substance is called an acid salt. It still con- 
tains hydrogen which is capable of forming an ion in solu- 
tion, and is therefore an acid; part of the hydrogen of the 
acid has been replaced by a metal, and the substance formed 
is therefore a salt. 



Acids, Bases, Salts, Valence. iii 

In a similar way, part of the hydroxyl groups of a 
base may be neutralized by an acid, and part remain un- 
changed. Thus we could have the reaction 

Ca(0H)2 + HCl =::: Ca(OH)Cl + H^O 
Such a substance is called a basic salt. 

When all of the hydrogen of the acid is replaced by a 
metal, the product is called a normal salt. Examples of 
these are sodium sulphate, Na2S04, potassium chloride, 
KCl, calcium nitrate, Ca(N03)2. Normal salts are in gen- 
eral neutral in reaction. A few, however, are acid while 
others are basic. 



CHAPTER XL ! 

COMPOUNDS OF NITROGEN. 

Introduction. As has been already stated, nitrogen 
constitutes a large portion of the atmosphere. The 
compounds of nitrogen, however, can not readily be ob- 
tained from this source, since nitrogen does not com- 
bine directly with many of the elements except at high 
temperatures. 

Combined with certain other elements, nitrogen occurs 
in the soil from which it is taken up by the plants and built 
into complex compounds composed chiefly of carbon, hy- 
drogen, oxygen and nitrogen. Animals feeding on these 
plants assimilate the nitrogen, so that this element is an 
essential constituent of both plants and animals. 

Decomposition of Organic Matter by Bacteria. All 
plants and animals then contain nitrogen as an essential ele- 
ment of their living matter. When the living matter dies, 
and undergoes decay, the carbon, hydrogen and oxygen are 
converted for the most part into carbon dioxide and water, 
while the nitrogen may appear either as free nitrogen or as 
ammonia, NHg, or as oxides of nitrogen. In the latter case, 
the oxides of nitrogen are subsequently converted into ni- 
trites or nitrates, which are respectively salts of nitrous and 
nitric acids. 

■ It has been shown that all such processes of decay are 
due to the action of different kinds of bacteria, each par- 
ticular kind effecting a definite change. The vast deposits 
of sodium nitrate which are found in warm, dry regions, 
especially in Chili and some parts of California, are thought 
to owe their origin to the decay of animal matter under the 
influence of these bacteria. 

(112) 



Compounds of Nitrogen. 113 

Decomposition of Organic Matter by Heat. When 
organic matter is strongly heated decomposition into simpler 
substances takes place, much as in the case of bacterial de- 
composition. If air is excluded, the nitrogen is liberated 
largely in the form of free nitrogen, although a considerable 
portion of it is combined with hydrogen to form ammonia. 
In the presence of plenty of air, oxides of nitrogen or nitric 
.acid are formed. When coal, which is a complex substance 
of vegetable origin consisting largely of carbon, but also 
<:ontaining hydrogen, oxygen and nitrogen, is heated in a 
-closed vessel, a part of the nitrogen present is changed into 
ammonia, and this is the chief source from which ammonia 
and its compounds are obtained. 

I. Compounds of Nitrogen with Hydrogen. — Am- 
monia. Several compounds consisting exclusively of 
nitrogen and hydrogen are known, but only one, ammonia, 
need be considered here. 
Preparation of Ammonia. 

a. Laboratory Method. In the laboratory ammonia is 
prepared from ammonium chloride obtained in the manu- 
facture of coal gas. When this substance is warmed with 
sodium hydroxide the reaction expressed in the following 
'equation takes place : 

XH.Cl + XaOH = XaCl -f NH.OH 

The ammonium hydroxide, NH^OH, so formed is un- 
stable, and breaks down into water and ammonia. 

XH.OH =XH3 + H,0 
The two equations can be combined into one as follows: 

XH.Cl + XaOH = XaCl + XH3 + H^O 

The cheaper calcium hydroxide, Ca(OH)2, is fre- 
quently used in place of sodium hydroxide, the equation 
ieing 

2 NH.Cl -f Ca(OH), = CaCU rf 2 H,0 + 2 XH3 



114 



Inorganic Chemistry. 



The ammonium chloride and calcium hydroxide are 
mixed together and placed in a flask arranged as shown in 
Fig. 32. The mixture is gently warmed, when ammonia gas 




Fig. 32 

is evolved, and being lighter than air, it is collected by up- 
ward displacement. ' 

Small amounts of ammonia can be made by passing elec- 
trical sparks through a mixture of hydrogen and nitrogen. The 
two elements combine directly. The action ceases however 
when a very small fraction of the two gases has combined, 
so that the method is not a practical one. 

b. Commercial Method. Nearly all the ammonia of 
commerce comes from the gas works. Ordinary illuminating; 
gas is made by heating coal in the absence of air. The pro- 
duct at first obtained contains a large number of compounds,, 
solids, liquids and gases. The solids and liquids are re- 
moved by cooling the product. Among the gaseous pro- 
ducts is ammonia. This is very soluble in cold water, while 
the gases valuable for illuminating purposes are not. The 
ammonia can therefore be separated from the other gases 
by passing the mixture through a vessel in which a spray 



Compounds of Nitrogen. 115 

of cold water plays upon the mixture, the ammonia being 
absorbed in the water. The liquid so obtained, called 
gas liquor, contains not only ammonia but other sol- 
uble substances. ]Most of these combine chemically with 
lime, while ammonia does not: if then lime is added to the 
gas liquor and the liquor is heated, the ammonia is driven 
out from it, being little soluble in hot water. It may be ab- 
sorbed again in pure cold water forming the aqua ammonia, 
or ammonia water of commerce. More frequently it is ab- 
sorbed in w^ater containing sulphuric or hydrochloric acids ; 
under these 'circumstances the ammonia combines with the 
acid to form respectively ammonium sulphate and ammo- 
nium chloride. These are solid compounds and can be 
obtained by evaporating the water and crystallizing out the 
salt. 

Physical Properties. Under ordinary conditions am- 
monia is a gas whose weight compared with that of an 
equal volume of air is 0.59 — that is, it is little more than 
half as heavy as air. It is easily condensed into a colorless 
liquid, and can now be purchased in liquid form in stout 
vessels. The gas is colorless and has a strong, suffocating 
odor. It is extrernely soluble in water, one liter of water 
at 0° and 760 mm. pressure dissolving 1,148 liters of the 
gas. In dissolving this large amount of gas the water ex- 
pands considerably, so that the density of the solution is 
less than that of water, the strongest solutions having a 
density of 0.88. The solution is called aqua ammonia. 

Chemical Properties. Ammonia will not support com- 
bustion, nor will it burn under ordinary conditions. In an 
atmosphere of oxygen it readily burns with a yellowish 
flame. When quite dry it is not a very active substance. 

Ammonium Hydroxide. The solution of ammonia in 
water is found to have strong basic properties. It turns red 
litmus blue ; it has a soapy feel ; it neutralizes acids form- 



ii6 Inorganic Chemistry. ^ 

ing salts with them. We have seen that a study of bases 
has led to the conclusion that their chief characteristic is 
that they contain one or more hydroxyl groups which act 
as ions in solution. It seems probable therefore that when 
ammonia dissolves in water it combines chemically with it 
according to the equation 

NH3 + H.O = NH4OH 
and that it is the substance NH4OH, called ammonium hy- 
droxide,, which has the basic properties, dissociating into 
the ions NH^ and OH. 

The group of atoms or radical NH^ plays the part of a 
metal and is called ammonium, since the names of metals 
generally have the ending -ium. The salts formed by the 
action of the base ammonium hydroxide on acids are called 
ammonium salts. Thus with hydrochloric acid there is 
formed ammonium chloride in accordance with the equation 

NH4OH + HCl = NH.Cl + H2O 

Similarly with nitric acid there is formed ammonium 
nitrate, NH4NO3, and with sulphuric acid, ammonium sul- 
phate, (NHJ2SO,. 

Ammonium hydroxide has never been obtained in a 
pure state. At every attempt to isolate it, the substance 
breaks up into water and ammonia. 

NH4OH = NH3 + H^O. 

It will be noticed that in the neutralization of ammo- 
nium hydroxide by acids the group NH^ replaces one hy- 
drogen atom of the acid, just as sodium does. The group 
therefore acts as a univalent metal. 

Volume Relations of Combining Gases. It has been 
found that when two volumes of ammonia are decomposed, 
three volumes of hydrogen are formed and one of nitrogen. 
If 'the two elements Vv^ere to combine directly therefore, one 
volume of nitrosfen would combine with three volumes of 



Compounds of Nitrogen. 117 

hydrogen to form two volumes of ammonia. In a former 
chapter we have seen that two volumes of hydrogen com- 
bine with one volume of oxygen to form two volumes of 
water vapor. The compound hydrochloric acid which 
has been frequently mentioned, can be made by the direct 
union of the gas chlorine and hydrogen, and it has been 
found that in this reaction one volume of chlorine combines 
with one volume of hydrogen to form two volumes of hydro- 
chloric acid gas. 

These results may be represented graphically thus : 

|"ir| + |"cr| = ihcii + iHcii 



H H + O = H,0 + H2O 



HHH-fN= NH3 + NH 



In the early part of the past century, Gay-Lussac, a 
distinguished French chemist, studied the volume relations 
of many combining gases, and concluded that similar rela- 
tions always hold. His observations are summed up in the 
following : 

Law of Gay-Lussac. When two gases combine chemi- 
cally, there is always a simple ratio between the volumes 
of the combining gases, and between the volume of either 
one of them and that of the product, provided it is a gas. 
By a simple ratio is of course meant the ratio of small whole 
numbers as 1:2, 2:3. 

II. Compounds of Nitrogen with Oxygen and Hy- 
drogen. Nitrogen forms several compounds with hydro- 
gen and oxygen, of which nitric acid, HNO3, and nitrous 
acid, HNO2, are the most familiar. 

a. Nitric Acid. Nitric acid is not found to any 
extent in nature, biit some of its salts, especially sodium ni- 
trate, NaNOg, and potassium nitrate, KNO3, are found in 
large quantities. From these salts nitric acid can be ob- 
tained. 



IIJ 



Inorganic Chemistry. 




Fig. 33 



Preparation of Nitric Acid. When sodium nitrate is 
treated with strong, cold sulphuric acid, no chemical action 
seems to take place. If however the mixture is heated in 
a retort, nitric acid is given off as a 
vapor and may be easily condensed to a 
liquid, by passing the vapor into a tube 
surrounded by cold water as shown in 
Fig. 33. 

2 NaN03+H2S04=Na,S04+2 HNO3 
The reason that this reaction takes 
place is not that the sodium has a 
stronger affinity for the 
SO4 group than for the 
NO3 group ; for if sodium 
sulphate and nitric acid 
are brought together the re- 
action goes the other way 
to quite an extent. 

( 1 ) NaoSO^ + 2HNO3 = H2SO, + 2NaN03' 

A reaction of this kind is said to be reversible, that is it 
can go in either way, depending upon the conditions. Un- 
der ordinary conditions the reaction 

(2) 2 NaN03 + H,SO, = Na.SO, -{- 2 HNO3 
goes on till a part of the sodium nitrate is changed into sul- 
phate, and then apparently ceases, the change in the one 
way being at the same rate as the change in the other. 
There is thus a kind of balance or equilibrium set up be- 
tween the tendency to go as in equation (i) and as in (2). 

If, when a state of balance has been reached, the 
amount of one of the four substances is changed, the equi- 
librium is disturbed. Thus if the nitric acid escapes from 
the mixture, more nitric acid will be formed to take its 
place. If the temperature rises above the boiling point of 
nitric acid, the nitric acid escapes as ' fast as formed and 



Compounds of Nitrogen. 119 

the reaction will go on until all the sodium nitrate has been 
"decornposed. 

Law of Mass Action. In a reaction which is capable of 
going in either w^ay, that is in a reversible reaction, it is 
therefore the relative amounts of the various substances 
present which determines in a large measure to what extent 
the reaction will proceed, and when it w^ll reach an equi- 
librium. The amounts of the substances which are available 
for the reaction are called the active masses of the sub- 
stances. The general statement can therefore be made : 
The extent to which a reversible reaction will take place 
depends upon the relative active masses of the substances 
taking part in the reaction. This is called the law of mass 
action. 

When we desire to have the reaction between sodium 
nitrate and sulphuric acid go on until it is complete, we 
must work under such conditions that the active mass 
of the nitric acid keeps very small. This will be the 
•case when the temperature of the mixture is kept above the 
toiling point of nitric acid, for the acid will then boil away 
as fast as formed, and never accumulate in the mixture. 

To prepare an acid therefore we need only treat one of 
its salts with some other acid which has a higher boiling 
point, and heat the mixture (if necessary) above the boiling 
point of the lower boiling acid. 

Physical Properties of Nitric Acid. Pure nitric acid is 
a colorless liquid which boils at about 86° and has a density 
of 1.56. The strong acid of commerce contains about 68% 
of HNO3, the remainder being water. Such a mixture has a 
density of 1.4. The strong acid fumes somewhat in moist 
air, and has a sharp choking odor. 

, Chemical Properties, i. Acid Properties. As the name 
indicates, this substance is an acid, and has all the properties 



120 Inorganic Chemistry, 

of that class of substances. It changes blue litmus red, ha^ 
a sour taste in dilute solutions, and neutralizes bases forming- 
salts. With potassium hydroxide and sodium hydroxide the 
reactions are represented by the equations : 

HNO3 + KOH = KNO3 + H2O 
HNO3 + NaOH =: NaNOg + H^O 
It also acts upon the oxides of most metals, forming a 
salt and .water as in equations : 

PbO + 2 HNO3 = PbCNOg)^ + H^O 
CuO + 2 HNO3 = Cu(N03)2 + H2O 
It is one of the strongest acids, by which is meant that it is 
largely broken up into its ions H and NO3 in solution. 

2. Decomposition on Heating. When boiled, or ex- 
posed for some time to sunlight, it suffers a partial decom- 
position according to the equation 

2 HNO3 = H2O + 2 NOo + O 
The substance NOo, called nitrogen peroxide, is a brownish 
gas which is readily soluble in water and in nitric acid. It 
therefore dissolves in the undecomposed acid, and imparts 
a yellowish color to it. Strong nitric acid highly charged 
with this substance is called fuming nitric acid. 

3. Nitric Acid as an Oxidizing Agent. As is shown 
by its formula, nitric acid contains a large amount of oxy- 
gen, and the reaction just mentioned shows that the com- 
pound is not a very stable one, easily undergoing decompo- 
sition. These properties should make it a good oxidizing 
agent, and we find that this is the case. Under ordinary 
circumstances, when acting as an oxidizing agent, it is de- 
composed according to the equation 

2 HNO3 = H^O + 2 NO + 3 O 
the oxygen being taken up by the substance oxidized, and 
not set free as is indicated in the equation. Thus if sulphur 



Compounds of Nitrogen. 121 

is oxidized by nitric acid, the oxygen will combine with 
sulphur, forming the substance SO3 and we will have 

3 O + S = SO3 
The two equations combined will be 

2 HNO3 + S = H2O + 2 NO + SO3 

Copper is oxidized by nitric acid, which converts it 
first of all into copper oxide. 

Cu + O = CuO. 
But copper oxide is very easily acted on by nitric acid^ 
as are most metallic oxides, forming a nitrate. 

CuO + 2 HNO3 = Cu(N03)2 + H2O 
So the complete action of nitric acid upon copper may 
be represented by the three equations : 

(i) 2 HN03=:H20 + 2 NO + 3 O 

(2) 3 + 3 Cu = 3 CuO 

(3) 3 CuO + 6 HNO3 = 3 Cu(N03). +3H,0 ^ 
These three equations may be combined into one by simply 
adding them together and canceling the equal quantities on 
each side the equation. 

8HNO3 + 3CU = 3Cu(N03)o + 2NO + 4H2O. 
This reaction is not peculiar to copper, but represents the 
usual action of nitric acid upon metals. The stages in it 
should be carefully noted. 

Salts of Nitric Acid — Nitrates. The salts of nitric 
acid are called nitrates. When heated to a high temperature 
nitrates undergo decomposition. In a few cases part uf the 
oxygen is given of¥ and a nitrite is left behind. 
' NaNOs = NaNO^ + O 

In other cases the decomposition goes further, and the 
metal is left as oxide 

Cu(N03)2 = CuO + 2 NO2 + O. 

All nitrates are soluble in water. 

b. Nitrous Acid. It is not difficult to obtain sodium 
nitrite, as the reaction given above indicates. Instead of 



\22 Inorganic Chemistry. 

merely heating the nitrate, it is better to heat it together 
with lead, when the reaction 

NaNOg + Pb = PbO + NaNO^ 
easily takes place. 

When sodium nitrite is treated with an acid such as 
sulphuric acid it is decomposed and nitrous acid is set free : 
2 NaNO^ + H,SO, = Na^SO, + 2 HNO^ 

The acid is very unstable however, and breaks up read- 
ily into water and an oxide of nitrogen : 
2 HNO, = H^O + N.O3 

Dilute solutions of the acid can, however, be obtamed, 

III. Compounds of Nitrogen with Oxygen. Nitro- 
gen combines with oxygen to form five different oxides. 
The formulas and names of these are as follows : 

N2O, nitrous oxide. 

NO, nitric oxide. 

N2O3, nitrogen trioxide or nitrous anhydride. 

NO2, nitrogen peroxide. 

N0O5, nitrogen pentoxide or nitric anhydride. 
These will now be briefly discussed. 

I. Nitrous Oxide — Laughing Gas. Ammonium ni- 
trate, like all nitrates, when heated undergoes decomposition ; 
but owing to the fact that it contains no metal the reaction 
is a peculiar one. 

NH.NOg = 2 H2O + N2O 

The oxide of nitrogen so form.ed is called nitrous oxide 
or laughing gas. 

Properties. It is a colorless gas having a slight odor. 
It is somewhat soluble in water, and in solution has a sHghtly 
sweetish taste. It is easily converted into a liquid and can be 
purchased in this form. When inhaled it produces a kind 
of hysteria, (hence the name "laughing gas") and in larger 
amounts unconsciousness and insensibility to pain. It has 



Compounds of Nitrogen. 123 

long been used as an anaesthetic for minor surgical opera- 
tions such as those of dentistry, but owing to its unpleasant 
after effects it is not so much used as it formerly was. 

Chemically, nitrotis oxide is remarkable for the fact that 
it is a very energetic oxidizing agent. Substances such as 
carbon, sulphur, iron and phosphorus burn in it almost as 
"brilliantly as in oxygen, forming oxides and setting free 
nitrogen. Evidently the oxygen in nitrous oxide cannot be 
held in very firm combination by the nitrogen. 

2. Nitric Oxide. AVe have seen that when nitric acid 
acts upon metals such as copper, the reaction represented 
"by the following equation takes place: 

3 Cu + 8 HNO3 := Cu(N03)2 + 2 NO + 4 H2O 

Nitric oxide, NO, is most conveniently prepared in this 

way. The metal is placed in the flask A (Fig. 34) and the 

acid added slowly through the funnel tube B. The gas es- 

capes through C and is collected over water. 

\) Pure nitric oxide is a colorless gas, slightly 

heavier than air, and is practically insoluble in 

^ ^ (^ water. It is a difficult gas to liquefy. Unlike 

yliiir nitrous oxide, nitric oxide does not part with 

fc=^ its oxygen easily, and burning substances in- 

troduced into this gas are usually extinguished. 

A few substances like phosphorus which have 

a very strong affinity for oxygen and which 

are burning energetically in the air, will con- 



Fig^. 34 tinue to burn in an atmosphere of nitric ox- 

ide. In this case the nitric oxide loses all of 
its oxygen and the nitrogen is set free as gas. 

It seems strange that nitrous oxide, though containing so 
much less oxygen than nitric oxide should give it up so much 
more readily. The explanation of this conduct seems to be 
that the molecules of nitrous oxide consist of a more unstable 
grouping of atoms than in the case of nitric oxide. The stability 
of a molecule seems to depend partly upon the character and 



124 Inorganic Chemistry. 

number of atoms in it, and partly upon the way in which they 
are arranged. So even with the same kind of atoms we may 
have one arrangement which is stable and one which is easily 
decomposed. 

When nitric oxide comes into contact with oxygen or 
with the air, it at once combines with oxygen even at ordi- 
nary temperatures, forming a reddish yellow gas of the 
formula NOg which is called nitrogen peroxide. This action 
is not energetic enough to produce a flame, but considerable 
heat is set free in it. Owing to the ease with which this 
action takes place it is impossible to say whether nitric oxide 
•has a taste or odor of its own, for what we taste or smell 
will not be nitric oxide, but nitrogen peroxide. 

3. Nitrogen Peroxide, NO2. This gas as we have just 
seen is prepared by allowing nitric oxide to come in contact 
with oxygen. It can also be made by heating a nitrate such 
as lead nitrate : 

Pb(N03)2 == PbO + 2 NO2 + O 
It is a reddish yellow gas of unpleasant odor, which is 
quite poisonous when inhaled. It is heavier than air and is 
easily condensed to a liquid. It dissolves in water, but this 
solution is not a mere physical solution ; the nitrogen per- 
oxide is decomposed, forming a mixture of nitric and nitrous 
acids. 

2 NO2 + HgO = HNO2 + HNO3 

Nitrogen peroxide will not combine with more oxygen; 
it will however give up a part of its oxygen to burning sub- 
stances, acting as an oxidizing agent. 

NO, = NO + O 

In this case the nitrogen peroxide is usually reduced to 
nitric oxide, which on exposure to the air will change once 
more to nitrogen peroxide. This substance can therefore be 
used to transfer the oxygen from the air to substances which 



Compounds of Nitrogen. 125 

will not burn in the air, but will burn in nitrogen peroxide. 
We shall see that a large industry depends upon this reaction. 
Anhydrides. The oxides NoOg or nitrogen trioxide, and 
JSFgOg or nitrogen pentoxide, bear a very interesting rela- 
tion to the acids of nitrogen. When dissolved in water, they 
combine with the water forming acids. 

N2O5 + H2O = 2 HNO3 

N2O3 + H2O = 2 HNO2 
On the other hand nitrous acid very easily decomposes 
yielding water and nitrogen trioxide, and by suitable means 
nitric acid likewise may be decomposed into water and 
nitric acid. 

2 HNO, = H,0 + N.O3 

2 HNO3 = H^O + N2O5 
In view of the close relation between these oxides and 
the corresponding acids they are called anhydrides of the 
acids ; N^03 being nitrous anhydride and N2O5, nitric 
anhydride. In general any oxide which will combine with 
water to form an acid, or which, together with water, is 
formed by the decomposition of an acid, is called an anhy- 
dride of that acid. We shall meet with many examples of 
such oxides. 



CHAPTER XII. 
SULPHUR. 

Occurrence. The element sulphur has been known 
from the earliest times, since it is widely distributed m 
nature and occurs in large quantities in the elementary form, 
especially in the neighborhood of volcanoes. Sicily has long 
been famous for its sulphur mines, and smaller deposits are 
found in other localities in Italy, Iceland, Mexico and 
Louisiana. In comibination, sulphur occurs in the form of 
sulphides, especially as sulphide of iron, copper, zinc and 
lead ; it also occurs as sulphates, particularly as calcium sul- 
phate or gypsum, CaSO^, 2 H2O and as barium sulphate or 
heavy spar, BaSO^. In smaller amounts it is found in a 
great variety of mineral forms, and is a constituent of many 
animal and vegetable substances. 

Extraction of Sulphur. The separation of crude sul- 
phur from the rock and earthy materials with which it is 
mixed is a very simple process. The ore from the mines 
is merely heated until the sulphur melts and drains away 
from the earthy impurities. The crude sulphur obtained in 
this way is purified by distillation. It is heated to boiling 
in a retort-shaped vessel made of iron, the exit tube of 
which opens into a cooling chamber of brickwork. When 
the sulphur vapor first enters the cooHng chamber it con- 
denses upon the walls as a fine crystalline powder called 
flowers of sulphur. As the condensing chamber becomes 
w^arm, the sulphur collects as a liquid in it, and is drawn off 
into cylindrical moulds, the product being called roll sulphur 
or brimstone. 

Physical Properties. Roll sulphur is g. pale yellow, 
crystalline solid without marked taste and with but a faint 

(126) 



Sulphur. 127 

odor. It is insoluble in water but is freely soluble in a few 
liquids as for example m carbon disulphide. 

From solution in these solvents it may be obtained in 
crystals which have the same color as roll sulphur and, like 
the latter substance, melt at 114.8°. Just above the melting 
point, sulphur forms a rather thin, straw-colored liquid ; as 
the temperature is raised, this liquid turns darker in color 
and becomes thicker until at about 235° it is almost black and 
so thick that the vessel containing it can be inverted with- 
out danger of the liquid running out. At higher tempera- 
tures it becomes thin once more, the color remaining dark, 
and boils at 448°, forming a yellowish vapor. On cooling 
the same changes take place in reverse order. 

When melted sulphur is allowed to cool until a part of 
the liquid has solidified, and the remaining liquid is then 
poured off, it is found that the solid sulphur remaining in 
the vessel has assumed the form of fine needle-shaped crys- 
tals which differ much in appearance from those which are 
obtained by crv^stallizing sulphur from solutions, and which 
are rhombic in form. The needle-shaped form is called 
monoclinic sulphur. The two varieties diiTer also in density 
and in melting point; the monoclinic sulphur melting at 120°. 
In a few days the needles become opaque and crumble into 
fragments which are found to have the rhombic form, and 
to melt at 114.8° like the other variety. Heat is set free in 
this transformation. 

Crystallography. In order to understand the difference 
between these two kinds of crystals, it is necessary to know 
something about crystals in general and the forms which they 
may assume. An examination of a large number of crystals 
has shown that although they may differ much in geometric 
form, they can all be considered as modifications of a few 
simple plans. The best way to understand the relation of one 
crystal to another is to look upon every crystal as having its 
faces and angles arranged in definite fashion about certain 



T28 



Inorganic Chemistry. 



imaginary lines drawn through the crystal. These lines are 
called axes, and bear much the same relation to a crystal as 
do the axis and parallels of latitude and longitude to the 
earth and a geographical study of it. All crystals can be 
referred to one of six simple plans or systems which have 
their axes as shown in the following drawings. 



7^ 




J < 



Fig. 35 



Fig. 





Fig. 








^ K 



^^ 



Fig. 39 



Fig. 40 



The names and characteristics of these systems are as 
follows: 

1. Isometric or regular system (Fig. 35). Three equal 
axes all at right angles. 

2. Tetragonal system (Fig. 36). Two equal axes and 
one of different length, all at right angles to each other. 

3. Orthorombic system (Fig. 37). Three unequal axes, all 
at right angles to each other. 

4. Monoclinic system (Fig. 38). 
and a third at right angles to one 
the other. 

5. Triclinic system (Fig. 39). 
to each other. 



Two axes at right angle"), 
of these, but inclined to 

Three axes all inclined 



Sulphur, 



129 



6. Hexagonal system (Fig. 40). Three equal axes in the 
same plane intersecting at angles of 60°, and a fourth at right 
angles to all of these. 

Every crystal can be imagined to have its faces and angles 
arranged in definite way around one of these systems of axes. 
A cube for instance is referred to plan i, an axis ending in. 
the center of each face; while in a regular octohedron, an axis 
ends in each solid angle. These forms are shown in Fig, 35. 
It will be seen that both of these figures belong to the same 
system, though they are very different in appearance. In the 
same way many geometric forms may be derived from each 
of the systems, and the light lines about the axes in the draw- 
ings show two of the simplest forms of each of the systems. 

In general, a given substance always crystallizes in the 
same system, and two corresponding faces of each crystal of 
it always make the same angle with each other. A few sub- 
stances, of which sulphur is an example, crystallize in two 
different systems, and the crystals differ in such physical 
properties as melting point and density. Such substances are 
said to be dimorphus. 

Amorphous Sulphur. Sulphur also exists in a third 
physical form called amorphous sulphur. Flowers of sul- 
phur always contains some of this m.odiiica- 
tion, and when the crystalline part of the 
flowers is dissolved away by carbon disul- 
phide, the amorphous sulphur is left undis- 
solved as a nearly white powder, it bemg in- 
soluble in carbon disulphide. 

Plastic Sulphur. When highly heated 
liquid sulphur is poured into cold water it 
assumes a gummy, dough-like form 
which is quite elastic. This can be 
seen in a very striking manner by 
distilling sulphur from a small, 
short necked retort such as is repre- 
sented in Fig. 41 and allowing the 
liquid to run directly into water. 




Fig. 41 



130 Inorganic Chemistry. ^ 

In a few days it becomes quite brittle and passes over inta 
ordinary rhombic sulphur. 

AUotropic Forms. In the study of oxygen it was ex- 
plained that the gas ozone is an allotropic form of oxygen. 
The modifications of sulphur just described are good examples 
of allotropic forms of a solid element. These forms differ not only 
in appearance, but also in their energy content. The amount of 
heat liberated by burning a gram of rhombic sulphur, for example,, 
is different from that which is set free in the combustion of a 
gram of monoclinic sulphur. 

Chemical Properties of Sulphur. When any of the 
forms of sulphur are sufficiently heated in oxygen or in the 
air, the sulphur burns with a pale blue flame forming sul- 
phur dioxide, SOo. Small quantities of sulphur trioxide, 
SO3, may also be formed. Most metals when heated with 
sulphur, combine directly w4th it forming metallic sulphides. 
In some cases the action is so energetic that the mass be- 
comes incandescent as has been seen in the case of iron 
uniting with sulphur. This property recalls the action of 
oxygen upon metals, and in general the metals which com- 
bine readily with oxygen are apt to combine quite readily 
with sulphur. 

Large quantities of sulphur are used in the manufacture 
of gunpowder, matches, vulcanized rubber and pure sul- 
phuric acid ; smaller quantities are used in a great variety 
of ways. 

Compounds of Sulphur with Hydrogen. Hydrosul- 
phuric Acid, HgS. A gas having the composition ex- 
pressed by the formula H2S is found in the vapors issuing^ 
from volcanoes, and in solution in the waters of many 
springs. It is also formed when organic, matter containing 
sulphur undergoes decay, just as ammonia is formed under' 
similar circumstances from nitrogenous matter. This gas 
is commonly called hydrogen sulphide. 



Q 



ff= 



Sulphur. 131 

Preparation. Hydrosulphuric acid is prepared in the 
laboratory by treating some sulphide — usually iron sul- 
phide — with an acid. 

FeS + 2 HCl = FeCL + H.S 
A convenient apparatus is shown in Fig. 42. A few 
lumps of iron sulphide are placed in the bottle A and diluted 
acid is poured upon it in small quantities at a 
time through the funnel tube B, the gas es- 
caping through the tube C. Iron sulphide is 
■B ^ a salt of hydrosulphuric acid, and this reaction 
is therefore similar to the one which takes 
|L^ place when sulphuric acid acts upon a nitrate. 

( \ 2 NaNOg + H2SO4 = Na^SO, + 2 HNO3 

In both cases when the salt and an acid are 
brou2:ht .tos^ether, there is a tendency for the 
reaction to go on until a state of balance is 
^ — reached between all four of the substances rep- 
resented in the equation. This balance is con- 
stantly disturbed by the escape of the gaseous acid set free, 
so that the reaction goes on until all of the original salt has 
been decomposed. The two reactions differ in that the first 
one is complete at ordinary temperatures, while in the case of 
sulphuric acid acting upon sodium nitrate, the reacting sub- 
stances must be heated so as to secure a temperature at which 
nitric acid will be a gas. 

Physical Properties. Hydrosulphuric acid is a col- 
orless gas having a weak, disagreeable taste and an ex- 
ceedingly offensive odor. It is rather sparingly soluble in 
water at ordinary temperatures, about 3 volumes dissolv- 
ing in 100 of water. In boiling water it is not soluble at 
all. In pure form it acts as a violent poison and even when 
diluted largely^ with air produces poison symptoms such 
as headache, dizziness and nausea. It is a little heavier 
than air, having a density of 1.18. 



Fig. 42 



^32 Inorganic Chemistry. 

Chemical Properties. The elements in the hydro- 
sulphuric acid are not very firmly held together and in conse- 
quence the compound is easily decomposed. It readily burns 
in the air, both of the elements in it forming oxides. 
H,S + 3 O = H,0 + SO, 

When there is not enough oxygen for both the sulphur 
and the hydrogen, the hydrogen takes the oxygen and the 
sulphur is set free. 

HoS + O = H2O + S 

On account of the ease with which the compound 
breaks down, and the strong affinity of both sulphur and hy- 
drogen for oxygen, hydrosulphuric acid is a strong reduc- 
ing agent, taking the oxygen away from substances con- 
taining it. 

We have seen, that when water vapor is passed over 
hot iron, an oxide of iron is formed and hydrogen is set free. 
The equation is 

4 H,0 + 3 Fe =.. Fe30, + 8 H 

A very similar action takes place when hydrosulphuric 
acid is passed over hot iron : 

4 H2S + 3 Fe = FegS^ + 8 H 

Hydrosulphuric acid is a weak acid. In solution in 
water it turns blue litmus red and neutralizes bases forming 
salts called sulphides. Many of these are insoluble in water 
and in acids, so that when hydrosulphuric acid is passed into 
solutions of certain metallic salts an insoluble sulphide is 
precipitated : 

CuCL + H,S = CuS + 2 HCl 

In this case, as in all others where a salt and an acid 
are brought together in solution, there is a tendency to an 
inferchange of the metal between the acids until a state of 
balance is reached. But copper sulphide is insoluble in the 
liquid in which the reaction occurs, and consequently is pre- 



Sulphur. 133 

cipitated, so that the amount in solution, that is the active 
mass of the sulphide, is kept very small. As a result a balance 
can not be reached, and the reaction goes on until all the 
copper has been precipitated. 

Because of the fact that some metals are precipitated as 
sulphides in this way, while some are not under similar cir- 
cumstances, hydrosulphuric acid finds a practical use in the 
problem of separating the metals from each other in chemical 
analysis, and is much used in the laboratory for this purpose. 

Oxides of Sulphur. Sulphur forms two well known 
oxides ; sulphur dioxide, SO2, and sulphur trioxide, some- 
times called sulphuric anhydride, SO3. 

Sulphur Dioxide. SO2. This substance is the chief 
product of the combustion of sulphur in air or in oxygen : 

S + 2 O = SO_. 
It is also formed when substances containing sulphur are 
burned. Thus when metallic sulphides are heated in a cur- 
rent of air, both the metal and the sulphur are changed into 
oxides 

ZnS + 30 = ZnO + SOo 
2 FeSo + 110 = FcoO. + 4 SOo. 

It is not surprising therefore that this substance is one 
of the gases which issues abundantly from volcanoes. 

Laboratory Preparation. In the laboratory sulphur 
dioxide is usually made by one of two methods. 

I. When a sulphite, which is a salt of sulphurous acid 
H2SO3, is treated with an acid, sulphurous acid is set free, 
and being very unstable, decomposes into water and sulphur 
dioxide. These reactions are expressed in the equations 
Na^SOg + 2 HCl = 2 NaCl + H^SOg 

H2SO, =: H2O + SO2 

In this case we have two reversible reactions depending 
on each other. In the first reaction 

NaoSOg + 2 HCl = 2 NaCl + H2SO3 



134 Inorganic Chemistry. ^ 

we should expect a balanced state to result, for none of the 
four substances in the equation are insoluble or volatile 
when water is present to hold them in solution. But the 
active mass of the H2SO3 is constantly diminishing owing to 
the fact that it decomposes as represented in the equation 

H.SOg = H,0 + SO2 
and the SO2 being a volatile gas, escapes. No sta,ble bal- 
ance can therefore result, since the active mass of the sul- 
phurous acid is Constantly being diminished because of the 
escape of SO2. 

2. Under some circumstances strong sulphuric acid 
acts as an oxidizing agfent and is itself reduced to sulphur- 
ous acid, the sulphurous acid so formed decomposing as 
represented in the last equation. Thus when copper is heated 
with the strong sulphuric acid the reactions expressed in 
the following equations take place: 

Cu + H2SO4 = CuO + H2SO3 

H2SO3 = H2O + SO2 . 

CuO + H^sb, = CuSO, + H2O 
The three equations may be combined into one : 

Cu + 2 H2SO, = CuSO, + SO2 + 2 H2O 

Physical Properties. Sulphur dioxide is a colorless 
gas, which at ordinary temperatures is about 2.2 times as 
heavy as air. It has a peculiar irritating odor which is 
familiar to every one since it is perceived when a sulphur 
match is burning. The gas is very soluble in water, one 
volume of water dissolving about 80 of the gas at 0°. It 
is easily condensed to a colorless liquid and can be purchased 
in this condition stored in strong glass bottles, siich as the 
one represented in Fig. 43. 

Chemical Properties, i. When sulphur dioxide is 
disst)lved in water, the solution acquires acid properties, and 
is found to contain an acid of the formula H2SO3, called 



Sulphur. 



135 




Fig. 4c 



sulphurous acid. Like /nitrous acid, sul- 
phurous acid is very unstable, and cannot be 
isolated in pure form ; on warming a solu- 
tion containing it, the acid is decomposed 
into water and the anhydride, as is the case 
with most acids whose anhydrides are gases. 
Salts of sulphurous acid can readily be made 
by treating a solution of the acid with a so- 
lution containing a soluble base, 
2 NaOH -j- H.SOg = Na^SOg + 2 H^O 
2. Solutions of sulphur dioxide act 
as good reducing agents ; that is sulpurous 
acid has the power to take up oxygen from 

the air or from substances rich in oxygen, and is by this 

reaction changed into sulphuric acid. 

H.SOg + O = H,SO, 
H.SOg + H,0, = H.SO^ + H.O 

The sulphites, i. e., the salts of sulphurous acid, like- 
wise have the power to take up oxygen, and are thereby 
changed into the corresponding sulphates. For this reason 
commercial sulphites are often contaminated with sulphates. 

3. Moist sulphur dioxide has strong bleaching prop- 
erties, acting upon many colored substances in such a way 
as to destroy their color. This is in part due to the reducing 
properties of the sulphurous acid. The action for the most 
part consists in partially reducing the color substance, and 
in many cases it happens that the substance into which it is 
changed is colorless. The gas is much used to bleach paper, 
straw goods and even such foods as canned corn. 

4. Sulphur dioxide also possesses marked antiseptic 
properties, and has the power to arrest fermentation. It is 
therefore used as a preservative. Enormous quantities are 
used in the manufacture of sulphuric acid. 



136 



Inorganic Chemistry. 



Sulphur Trioxide. .When sulphur dioxide and oxy- 
gen are heated together at a rather high temperature a 
small amount of sulphur trioxide, SO3, is formed, but the 
reaction is slow and incomplete. If however the heating 
takes place in the presence of very fine platinum dust the 
reaction is rapid and nearly complete. 

The experiment can be easily performed by the use of 
the apparatus shown in Fig. 44, the fine platinum being se- 
cured by moistening asbestos fiber with a solution of plati- 






\, 



Jg^ 






Fig. 44. 

num chloride and igniting it in a flame. The fiber, . covered 
with fine platinum, is placed in- a tube of hard glass which is 
then heated with a burner, while sulphur dioxide and air are 
passed into the tube. Union takes place at once, and the 
strongly fuming sulphur trioxide escapes from the jet at the 
end of the tube, and may be condensed by surrounding the 
receiving tube with a freezing mixture. 

Catalysis.. It has been found that many chemical re- 
actions are, like this one, much influenced by the presence of 
substances which do not themselves seem to take a part in the 
reaction, and are left apparently unchanged after it has ceased. 
In most cases these reactions are ones which go on very 
slowly under ordinary circumstances, but are much hastened 
by the presence of the foreign substance. Substances which bring 
about reactions in this way, or hasten very slow reactions, are 
said to act as catalytic agents, and the whole action is called 
catalysis. Just how the action is brought about is not well 
understood. We have already had several instances of such 
action. Oxygen and hydrogen combine with each other at 
ordinary temperatures in the presence of platinum sponge^ 



Sulphur. 137 

though their ordinary kindling temperature is rather high. 
Potassium chlorate, when heated with manganese dioxide, 
gives up its oxygen at a much lower temperature than when 
heated alone. Manganese dioxide dropped into a strong so- 
lution of hydrogen dioxide causes it to decompose very rap- 
idly into water and oxygen. On the other hand the catalytic 
agent sometimes retards chemical action. A solution of hydro- 
gen dioxide decomposes more slowly when it contains a little 
phosphoric acid, than when perfectly pure. 

Properties. Sulphur trioxide is a colorless liquid 
which solidifies at about 15° and boils at 46°. A trace of 
moisture causes it to solidify into a mass of silky white 
crystals, somewhat resembling asbestos fiber in appearance. 
In contact with the air it fumes strongly, and thrown upon 
water it dissolves with a hissing sound and the liberation 
of a great deal of heat. The product of this reaction is sul- 
phuric acid, so that sulphur trioxide is an anhydride. 

SO3 + H^O == H3SO4 

Manufacture of Sulphuric Acid. 

I. Contact Method. Advantage is taken of the reac- 
tions just described, in the preparation of sulphuric acid on 
the large scale. Large iron tubes are packed with some porous 
material, such as calcium and magnesium sulphates, which 
contains a suitable catalytic substance scattered through it. 
The catalizers most used are platinum powder and iron ox- 
ide. Purified sulphur dioxide and air are passed through the 
tubes which are kept at a temperature of about 400°. Sul- 
phur trioxide is formed, and as it issues from the tube it is 
absorbed in water or dilute sulphuric acid. The process is 
continued until all the water in the absorbing vessel has been 
changed into sulphuric acid, so that a very concentrated acid 
is made in this way. Sulphuric acid has the power to dis- 
solve sulphur trioxide, and such a solution is called fuming 
sulphuric acid. 



138 Inorganic Chemistry. 

"2. Chamber Process. The method of manufacture ex 
clusively used until recent years, and still in very extensive 
use, is much more complicated, but it is exceedingly inter- 
esting as showing the way in which difficult chemical prob- 
lems can be solved. As has been stated, sulphur dioxide 
combines very slowly with oxygen to form sulphur triox- 
ide, and this is the practical difficulty in the manufacture of 
sulphuric acid. Indirectly this combination can be brought 
about in the following way. 

Sulphur dioxide is conducted into a very large lead 
hned chamber containing the gas NO2 ; steam and air are 
at the same time blown into the chamber. Reaction takes 
place according to the equation 

2 SO, + 2 NO2 + H2O -f O =; 2 SO^COH) (NO2). 

The product formed in this reaction is called nitro- sul- 
phuric acid. This substance is immediately acted upon by 
an additional amount of steam, sulphur dioxide and air, 
decomposition taking place according to the equation 
SO2+2 S02(OH)(N02)+2 H2O+2 0=3 H,SO,-f2 NO2 

The nitrogen peroxide formed in this last reaction can 
now enter into combination with a new amount of sulphur 
dioxide, steam and oxygen and the series of reactions go on 
indefinitely. It takes a great deal of careful watching to 
keep the substances present in just the right proportions. 
Although the equations do not indicate it, there is a constant 
loss of nitrogen peroxide through leakage, and through its 
solubility in the liquids formed; hence it, must be replaced 
from time to time. 

The nitrogen peroxide is obtained by the action of sul- 
phuric acid upon sodium nitrate, the nitric acid so obtained 
decomposing in the lead chamber to form the necessary oxide 
of nitrogen. Two or three parts of nitrate are used up for 
every 100 parts of sulphuric acid produced. The sulphur di- 
oxide is made by burning iron pyrite, FeS2, or is obtained as 



Sulphur. 139 

a bj^-product in the treatment of sulphide ores. Occasionally 
sulphur itself is burned to furnish it. 

The sulphuric acid so formed, together with the excess 
of condensed steam, collects upon the floor of the lead 
chamber, the acid liquid so obtained containing about 55% 
sulphuric acid. The product is called chamber acid, and is 
quite impure ; but for many purposes such as the manufac- 
ture of fertilizers it needs no further treatment. It can be 
concentrated by boiling it in vessels made of iron or platinum 
which resist the action of the acid, nearly all of the water 
iDoiling off and leaving the high boiling acid behind. Pure 
concentrated acid can best be made by the contact process 
Avhile the chamber process is cheaper for dilute, impure acid. 

Properties of Sulphuric Acid. Sulphuric acid is a 
colorless, oily liquid nearly twice as heavy as water. The 
ordinary concentrated acid contains about 2% of watei, has 
a density of 1.84 and boils at 338°. It is sometimes called 
oil of vitriol, since it was formerly made by distilling green 
vitriol. It possesses chemical properties which make it one 
of the most important of chemical substances. 

I. Action on Water. Sulphuric acid has a very 
great affinity for water, and is therefore an effective dehy- 
drating agent. Gases which have no chemical action upon 
sulphuric acid can be freed from water vapor by bubbling 
them through the strong acid. When the acid is diluted 
with water much heat is set free, and care must be taken 
to keep the liquid thoroughly stirred during the mixing, 
and to pour the acid into the water — never the reverse. 

Xot only can sulphuric acid absorb water, but it will 
often withdraw the elements hydrogen and oxygen from 
a compound containing them, decomposing the compound, 
and combining with the water so formed. For this reason 
most organic substances such as sugar, wood, cotton and 
wool fiber, and even flesh, all of which contain much oxy- 



140 Inorganic Chemistry. ^ 

gen and hydrogen, are charred or burned by the action of 
the strong acid. 

2. Action on Salts. We have repeatedly seen that an 
acid of high boihng point heated with the salt of some acid 
of lower boiling point will drive out the low boiling acid. 
The boiling point of sulphuric acid (338°) is higher than 
that of almost any common acid, so that this acid is much 
used to decompose salts in the preparation of acids from 
them. 

3. Oxidizing Action. Sulphuric acid contains a large 
percentage of oxygen and like most substances rich in oxy- 
gen it loses a part of it when heated with substances having 
a strong affinity for the element. When the strong acid is 
heated with carbon, sulphur, metals and many other sub- 
stances this kind of action takes place, as is shown in the 
equations following: 

C + 2 H2SO4 = CO2 + 2 H2SO3 
S + 2 H2SO4 = SO_. + 2 H2SO3 
Cu + H2SO4 = Cub + H2SO3 
In each case the sulphurous acid formed in the reaction 
breaks down into sulphur dioxide and water. When a 
metallic oxide is produced by the oxidizing power of sul- 
phuric acid, it dissolves in an additional amount of the acid 
forming a sulphate. Thus with copper the complete reac- 
tion is 

Cu + 2 H,SO, = CuSO^ + H2O + SO2 
In some cases sulphuric acid may lose all of its oxygen 
and be reduced to hydrosulphuric acid, H2S. With the sub- 
stance HI we have such a reaction under some conditions. 
H2SO4 + 8 HI = H2S + 4 H2O + 8 I 

4. Action as an Acid. In dilute solutions sulphuric 
acid acts as any other acid, neutralizing bases, forming salts 
with metallic oxides and liberating hydrogen when acted 
upon by metals such as zinc and iron. 



Sulphur. 141 

Salts of Sulphuric Acid — Sulphates. The salts of 
sulphuric acid are among the most familiar of all the salts, 
and many of them have familiar uses. Copperas, (iron 
sulphate), blue vitriol (copper sulphate), Epsom salt (mag- 
nesium sulphate) are good examples of this. Many sul- 
phates are important minerals, prominent among these being 
g-ypsum (calcium sulphate) and barytes (barium sulphate). 

Thiosulphuric Acid — Thiosulphates. Many other 
.^cids of sulphur containing oxygen are known, but none of 
them are of great importance. Most of them cannot be 
prepared in a pure state, and are known only through their 
salts. 

When sodium sulphite is boiled with sulphur the two 
substances combine forming a salt which has the composition 
represented in the formula Na2S203 to which the name so- 
dium thiosulphate has been given. 

NaoSOg + S = Na^S^Og 

This reaction is quite similar to the action of oxygen upon 
sulphites. 

Na.SOg + O = Na^SO, 

More commonly the salt is called sodium hyposulphite, or 
**hypo." It is extensively used in photography, and since 
it has the power to absorb large amounts of chlorine, it is 
used in the bleaching industry to remove the chlorine left 
upon cotton cloth after it has been bleached by chlorine. 
It is also used as a disinfecting agent. 

Monobasic and Dibasic Acids. Such acids as hydro- 
chloric acid, which can yield only one hydrogen ion, are 
called monobasic acids, while those giving more than one 
are called di-, tri-, or tetrabasic acids, according as they give 
two, three of four hydrogen ions. The three acids of sul- 
phur are dibasic acids. It is therefore possible for them 
to form both normal and acid salts. 



142 Inorganic Chemistry. ^ 

The acid salts can be made in two ways : ti.e acid may 
be treated with only half enough base to neutralize it : 

NaOH + H^SO, = NaHSO, + H^O 
or a normal salt rnay be treated with the free acid : 

Na.SO^ + H2SO4 = 2 NaHSO^ 

Acid sulphites and sulphides may be made in the same 
ways. 

Carbon Bisulphide. CS2. When sulphur vapor is 
passed over highly heated carbon, the two elements com- 
bine forming carbon disulphide, CS2, just as oxygen and car- 
bon unite to form carbon dioxide, CO2. The operation is 
carried out in such a way that the sulphur and carbon are 
heated together by means of a strong electrical current 
passed through the carbon, heating it to a high temperature. 
The substance is a heavy, colorless liquid, possessing when 
pure a pleasant ethereal odor. On standing for some time, 
especially when exposed to sunlight, it undergoes a slight 
decomposition and acquires a most disagreeable, rancid 
odor. It has the power of dissolving many substances such 
as gums, resins and waxes which are not soluble in most 
liquids, and its chief use is as a solvent for such substances. 

It boils at a low temperature (46°) and its vapor is very 
inflammable, burning in the air to form carbon dioxide and 
sulphur dioxide. 

CS2 + 6 = C02 + 2 SO2 

Comparison of Sulphur and Oxygen. A comparison 
of the formulas and the chemical properties of corresponding 
compounds of oxygen and sulphur brings to light many 
striking similarities. The conduct of H2S and HgO toward 
many substances has been seen to be very similar ; the ox- 
ides and sulphides of the metals have similar formulas and 
undergo many similar reactions. Carbon dioxide and di- 
sulphide are prepared in similar ways and have many analo- 
gous reactions. It is clear therefore that these two elements 



Sulphur. 143 

are far more closely related to each other than to any of 
the other elements so far studied. 

Selenium and Tellurium. These two very uncommon 
elements are still more closely related to sulphur than is oxygen. 
They occur in very small quantities and are usually found as- 
sociated with sulphur and sulphides, either as the free elements, 
or more commonly in combination with metals. They form 
compounds with hydrogen of the formulas H2Se and H2Te; 
these bodies are gases with properties very similar to those 
of H2S. They also form oxides and oxygen acids which re- 
semble the corresponding sulphur compounds. The elements 
even have allotropic forms corresponding very closely to those 
of sulphur, selenium corresponding more perfectly, while tel- 
lurium is in appearance more metallic. 

Telhirium is sometimes found in combination with gold 
and copper, and occasions some difficulties in the refining of 
these metals. The elements have very few practical applications. 



CHAPTER XIII. 



PERIODIC LAW. 



A number of the elements have now been studied some- 
what closely. The first three of these, oxygen, hydrogen and 
nitrogen, while having some physical properties in common 
with each other, have almost no point of similarity as re-" 
gards their chemical conduct. We have seen that oxygen 
and sulphur, on the other hand, while quite different 
physically have much in common in their chemical prop- 
erties. 

About 80 elements are now known, and if all of these 
should have properties as diverse as do oxygen, hydrogen 
and nitrogen, the study of chemistry would plainly be a very 
difficult and complicated one. If, however, the elements can 
be classified in groups the members of which have very simi- 
lar properties, the study will be a much more simple matter. 

Natural Grouping of Elements. Efforts have been 
made for a long time to find some natural principle in 
accordance with which the elements can be satisfactorily 
classified. Of course with such a large number of elements, 
many chance resemblances will necessarily occur. But it 
was felt that there must be some real cause in the nature of 
the elements sulphur and oxygen which occasions their simi- 
larity ; and that the same cause should group other elements 
together in families of similar properties. 

Metals and Non-Metals. The division of the ele- 
ments which most readily suggests itself, and which was first 
made, was into metals and non-metals. This division was 
founded largely upon physical properties, a metal being an 
element of high density, bright luster, opaque towards light, 
a good conductor of electricity, at least somewhat malleable 

(144) 



Periodic Law. 145 

and ductile ; and since elements possessing these properties 
were usually base forming in character, the ability to form 
iDases came to be regarded as a characteristic property of 
metals. The non-metals were acid forming with physical 
properties opposite to those of the metals. 

This division is unsatisfactory for two reasons. There 
are so many elements in each group that little is gained by a 
division. It is not based on a fundamental property, and the 
properties of some of the elements make it difficult to as- 
sign them to either of the divisions. Potassium is undoubt- 
edly a metal, yet it is one of the lightest of solids. Iodine, a 
typical non-metal, has a high metallic luster. Many ele- 
ments can form either an acid or a base depending on the 
circumstances. 

Triad Families. It was pointed out by Doebereiner 
in 1825 that there is an interesting relation existing between 
the atomic weights of chemically similar elements. Lithium, 
sodium and potassium resemble each other in much the 
same way as do sulphur and oxygen. The atomic weight of 
sodium is almost an arithmetical mean between those of the 

7 -I- 39 
Other two — ^ — —23. In many chemical and physical prop- 
erties sodium stands midway between the other two. Ba- 
rium, strontium and calcium form a similar group, with 
■strontium a mean between the others, both in atomic weight, 
and in most other properties. A number of other triad 
groups have been found. But among 80 elements whose 
atomic weights range all the way from i to 240, such agree- 
ments might be mere chance, and many such arithmetical 
groups can be found where there is no chemical similarity. 
Periodic Law. In 1869 the Russian chemist Men- 
•delejeff devised an arrangement of the elements based on 
their atomic weights, which has proved to be of great ser- 
A^ice in the comparative study of the elements. A few 



146 Inorganic Chemistry. 

months later, the German, Lothar Meyer, independently- 
suggested the same ideas. Since this arrangement has 
brought to light a great generalization in respect to the ele- 
ments, the generalization is known as a law — the periodic 
law. The arrangement suggested by Mendelejeff, modified/ 
somewhat by other workers and niore recent investigations, 
is as follows. 

Arrangement of the Periodic Table. Beginning 
with lithium which has an atomic weight of 7, the elements 
are arranged in a horizontal row in order of their atomic 
weights. Thus : 

Li (7), Gl (9), B (10.9), C (11.9), N (13.9), O (15.8), F (18.9) 

These seven elements all differ markedly from each other. 
The eighth element, sodium, is very similar to lithium. It is 
placed just under lithium and a new row follows : 

Na (22.9). Mg (24.2), Al (26.9), Si (28.2), P (30.8), S (31.8), CI (35.2) 

When the sixteenth element, potassium, is reached, it is 
placed under sodium, to which it is very similar, and serves 
to begin a third row : 

K (38.9), Ca (39.8), Sc (43.8), Ti (47.7), V (50.8), Cr (51.7), Mn (-54.6). 

Not only is there a strong similarity between sodium, lithium 
and potassium, which have been placed in a vertical row be- 
cause of this resemblance to each other, but the elements in 
the other vertical rows exhibit -much of the same kind of 
similarity among themselves, and evidently form little nat- 
ural groups. 

The three elements following manganese, namely, iron^ 
nickel and cobalt, have atomic weights near together, and are 
very similar chemically. They do not strongly resemble any 
of the elements so far considered, and are accordingly placed 
in a group by themselves, following manganese. A ncvv 
row is begun with copper, which somewhat resembles the 
elements of the first vertical column. Following the fifth 
and seventh rows are groups of three closely related ele- 



Periodic Law. 147 

ments so that the completed arrangement has the appear- 
ance represented in the table on page 148. 

A single horizontal row is called a short period. A 
group of elements in the eighth column, together with the 
short period just before the three, and the one just after — 17 
elements in all — is called a long period. A short period just 
preceding a group of three in the eighth column is called 
an even period; one just following is called an odd period. 
It will be noticed that the elements of the even periods are 
placed at one side (at the left) of the vertical column in 
which they fall while the elements of the odd periods are 
placed at the other side (at the right). This divides each 
group into two families. Arranging the elements in this 
way, certain very striking facts appear. 

Properties of Elements Vary with Atomic Weight. 
There is evidently a close relation between the properties o£ 
an element and its atomic weight. Lithium, at ^e begin- 
ning of the first period, is a very strong base-forming ele- 
ment, with pronounced metallic properties. Glucinum, fol- 
lowing lithium, is less metallic, while boron has some me- 
tallic and some acid-forming properties. In carbon, all me- 
tallic properties have disappeared, and the acid forming 
properties are more marked than in boron. These become 
still more emphasized as we pass through nitrogen and oxy- 
gen, until on reaching fluorine, we have one of the strongest 
acid-forming elements. The properties of these seven ele- 
ments therefore vary regulary with their atomic weights, or 
in mathematical language, are functions of them. 

Periodic Variation of Properties. The properties of 
tha first seven elements vary continuously — that is steadily 
— away from metallic, and toward acid-forming properties. 
If lithium were the lightest of all elements in atomic weight, 
and fluorine the heaviest, so that in passing from one to the 
other we had gone over all the elements, we could say 



148 



Inorganic Chemistry. 



Group 
VIII. 






<OC0«3 




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Group 
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Periodic Law. ■ 149 

that the properties of elements are a continuous func- 
tion of their atomic weights. But fluorine is an element 
of small atomic weight, and the one following it, sodium, 
breaks the regular order, for in it reappear all the character- 
istic properties of lithium. Magnesium, following sodium, 
bears much the same relation to glucinum that sodium does 
to lithium, and the properties of the elements in the second 
row keep varying much as in the first row, until potassium 
is reached, when another repetition begins. The properties 
of the elements do not vary continuously, therefore, with 
atomic weight, but at regular intervals there is a repetition, 
or period. The periodic law may be thus stated : The 
properties of an element are a periodic function of its atomic 
weight. 

The two Families in a Group. While all the ele- 
ments in a given vertical column bear a general resemblance 
to each other, the ones belonging to an even .period, and 
placed to the left in a group column, are very strikingly 
similar to each other, and the same is true of the elements 
at the right, in the odd periods. Thus lithium, sodium and 
potassium are very much alike ; so too are copper, silver and 
gold. The resemblance between lithium and copper, or so- 
dium and silver, is much less marked. The periodic arrange- 
ment therefore divides the elements into a number of fam- 
ilies, each containing three or four members between which 
there is a great similarity. 

Family Resemblances. Let us now inquire more 
closely in what respects the elements of a family resemble 
each other. 

I. Valence. In general the valence of the elements in 
a family is the same, and the formulas of their compounds 
are therefore similar. If we know that the formula of sodium 
chloride is NaCl, it is pretty certain that the formula of po- 
tassium chloride will be KCl — not KCI2 or KCI3. The for- 



150 Inorganic Chemistry. 

mulas R2O, RO, etc., placed below the columns show the for- 
mulas of the oxides of the elements in the column, while the 
formulas RH, RH^, etc., show the composition of the com- 
pounds formed with hydrogen or chlorine. 

2. Chemical Properties. The chemical properties of 
the members of a family are quite similar. If one member 
is a metal, the others usually are : if one is a non-metal, so 
tQo are the others. The families in the first two columns 
consist of metals, while the elements found in the last two 
columns form acids. There is in addition a certain regularity 
in properties of the elements in one family. If the element 
at the head of the family is a strong acid-forming element, 
this property is likely to diminish gradually, as we pass to 
the members of the family with higher atomic weight. Thus 
phosphorus is strongly acid-forming, arsenic less so, anti- 
mony still less so, while bismuth has almost no acid forming 
properties. We shall meet with many illustrations of this fact. 

3. Physical Properties. In the same way, the physical 
properties of the members of a family are in general some- 
what similar, and show a regular gradation as we pass from 
element to element in the family. 

Value of the Periodic Law. It is at once evident 
that such regularities very much simplify the study of chem- 
istry. A thorough study of one element of a family makes 
the study of the other members a much easier task, since so 
many of the properties and chemical reactions of the elements 
are so similar. If an element with atomic weight 100 should 
be discovered, it would be possible to predict many of its 
chief characteristics from the position it would occupy in 
the periodic table. 

When the periodic law was first formulated, there were 
a number of gaps in it, as there still are, which evidently be- 
longed to elements at that time unknown. From their posi- 
tion in the table, Mendelejeff predicted with great minute- 



Periodic Law. 151 

ness, the properties of the elements which he felt sure would 
one day be discovered to fill these places. Three of them, 
scandium, germanium and gallium, were found within fif- 
teen years and their properties agreed in a remarkable way 
with the predictions of Mendelejeff. 

The physical constants of many of the elements did 
not at first agree with those demanded by the law, and a 
further study of many of such cases showed that errors 
had been made, so that the law has done much service in 
indicating probable error. 

Imperfections of the Law. There still remain a 
good many features . which must be regarded as imperfec- 
tions in the law. Most conspicuous is the fact that the ele- 
ment hydrogen has no place in the table. At first there 
seemed to be no place for argon, but when the other in- 
active elements were found, it was seen that they together 
formed a group just preceding the first group, as shown in 
the table. 

In some of the groups, elements appear in one of the 
families, while all of their properties show that they belong 
in the other. Thus sodium belongs with lithium and not 
with copper ; fluorine belongs with chlorine and not with 
manganese. There are two instances where the elements 
are transposed in order to make them fit into their proper 
group. Tellurium should precede iodine, and argon should 
precede potassium. The table separates some elements alto- 
gether, which in many respects have closely agreeing prop- 
erties. Iron, chromium and manganese are all in different 
groups, although they are similar in many respects. 

The system is therefore to be regarded as but a partial 
and imperfect expression of some very important and fun- 
damental relation between the substances which we know 
as elements, the exact nature of this relation being not as yet 
completely clear to us. 



CHAPTER XIV. 

CHLORINE FAMILY. 

General. The four elements fluorine, chlorine, bro- 
mine and iodine form a strongly marked family of ele- 
ments and illustrate very clearly the way in which the 
niembers of a family in the periodic group resemble each 
other and the character of the differences which we may 
expect to find between the individual members. The com- 
pounds of the last three elements are found extensively in 
sea salt, and on this account the name halogens, signifying: 
"producers of sea salt," is sometimes applied to this family. 
It is more frequently called the chlorine family. 

FLUORINE. 

Occurrence. The element fluorine occurs in nature 
most abundantly as the mineral fluorspar, CaF2, as cryolite, 
NagAlF^j and in the complex mineral apatite, Ca3(P04)2, 
CaF^. 

Preparation. All attempts, to isolate the element 
resulted in failure until recent years. Methods similar to 
those which succeed in the preparation of the other elements 
of the family cannot be used for the reason that the fluorine 
as soon as liberated, instantly combines with hydrogen of 
the water always present or with the materials of which the 
apparatus is made. The preparation of fluorine was finally 
accomplished by the French chemist Moissan in the follow- 
ing way. Perfectly dry hydrofluoric acid was condensed 
to a liquid and placed in a U-shaped tube made of platinum 
(or copper), which was furnished with electrodes and de- 
livery tubes as shown in Fig. 45. A current will not pass 
through this liquid any more than through pure water ; but 

(152) 



Chlorine Family. 



153 




Fig. 45. 



if an electrolyte such as potassium fluoride is added the cur- 
rent will then pass. Hydrogen is set free at one electrode 
and fluorine at the other. 

Properties. Fluor/rx is a gas 
of slightly yellowish color, which 
can be liquefied by cold and pres- 
sure, the liquid boiling at — 187^ 
under atmospheric pressure. It is 
extremely active chemically, being 
at ordinary temperatures the most 
active of any of the elements. 

It combines with all the com- 
mon elements save oxygen, very 
often with incandescence, and the 
liberation of much heat. It has a 
great affinity for hydrogen and is 
able to withdraw it from its com- 
pounds with other elements. Be- 
cause of its great activity it is extremely poisonous. 

Hydrofluoric Acid. Hydrofluoric acid is readily ob- 
tained from fluorspar by the action of strong sulphuric acid. 
The equation is 

CaF. +H2SO, = CaSO, + 2 HF 
In its physical properties it resem.bles the hydrogen acids of 
the other elements of this family, being however, more easily 
condensed to a liquid. The anhydrous acid boils at 19° and 
can therefore be prepared at ordinary pressures. It is freely 
soluble in water, and a strong solution — about 50% — is 
prepared, for the market. Its fumes are exceedingly irri- 
tating to the respiratory organs and several chemists have 
met their death by accidently breathing them. 

Chemical Properties. Hydrofluoric acid resembles 
hydrochloric acid in many particulars, only that it is more 
energetic, acting upon substances which are not attacked by 



154 Inorganic Chemistry. ^ 

the latter acid. It acts very vigorously upon organic matter 
even a drop of the strong acid making a sore on the skin 
which is very painful and slow in healing. Its most charac- 
teristic property is its action upon silicon dioxide SiOg, with 
which it forms the gas silicon tetrafluoride, SiF^, and water, 
as shown in the equation 

SiOo + 4 HF =: SiF, + 2 H^O 
Glass consists of certain compounds of silicon, which are 
likewise acted on by the acid so that it cannot be kept in 
glass bottles. It is preserved in flasks made of wax or gutta 
percha. 

Etching. Advantage is taken of this reaction in etching 
designs upon glass. The glass vessel is painted over with 
a protective paint upon which the acid will not act, the parts 
which it is desired to make opaque being unprotected. A 
mixture of fluorspar and sulphuric acid is then painted over 
the vessel and after a few minutes the vessel is washed clean. 
Wherever the hydrofluoric acid comes in contact with the 
glass, it acts upon it, destroying its luster and making it opaque, 
so that the exposed design will be etched upon the clear glass. 
Frosted glass globes are often made in this way. 

Fluorides. The fluorides have formulas similar to 
the chlorides, bromides and iodides, but their properties are 
usuallv very different. This is seen in the solubility of the 
salts. Those metals whose chlorides are insoluble form 
soluble fluorides, while many of the metals forming soluble 
chlorides form insoluble fluorides. Most of the soluble 
fluorides are poisonous. Fluorine does not form any oxides, 
or any acids containing oxygen. 

CHLORINE. 

Historical. While studying the action of hydro- 
chloric acid upon the mineral pyrolusite in 1774, Scheele, a 
Swedish chernist, obtained a yellowish gaseous substance to 
which he gave a name in keeping with the phlogiston theory 



Chlorine Family. 



155 



then current. It was supposed to be a compound contain-' 
ing oxygen until the EngHsh chemist, Sir Humphrey Davy, 
in 181 o proved it to be an element and renamed it chlorine, 
the term signifying yellowish green which is the color of 
the gas. 

Occurrence. Chlorine does not occur free in nature 
but its compounds are widely distributed. For the most part 
it occurs in combination with the metals in the form of 
chlorides, those of sodium, potassium and magnesium being 
the most abundant. Nearly all salt water contains these sub- 
stances, particularly sodium chloride, and very large salt 
beds of similar composition are found in many countries. 
Preparation. 

a. Laboratory Method. In the laboratory chlorine is 
made by warming the mineral pyrolusite, (manganese diox- 
ide, MnOa), with strong hydrochloric acid. The 
first reaction seems to be similar to the action of 
acids upon oxides in general. 

■ MnO, + 4HCI = MnCl, + 2H,0. 

The manganese compound so 
formed is very unstable however and 
breaks down according to the equation 
MnCl, = MnCU + 2 CI 
The manganese dioxide and the 
hydrochloric acid are brought together 
in a flask as represented in Fig. 46, 
and gentle heat is applied. The 
rate of evolution of the gas is regu- 
lated by the amount of heat ap- 
plied, and the gas is collected by 
downward displacement. As the 
equations show, only half of the chlorine present in the 
hydrochloric acid is liberated. Instead of using hydro- 
chloric acid in the preparation of chlorine, it will serve just 




Fig. 46. 



156 Inorganic Chemistry. 

as well to use a mixture of sodium chloride and sulphuric 
acid, since these two react to form hydrochloric acid. The 
following equations then express the changes. 

4 NaCl + 2 H2SO4 = 4 HCl + 2 Na^SO, 
MnO^ + 4HCI = MnCl^ + CU + 2H2O 
The manganese chloride is further acted on by the' sul- 
phuric acid present as follows : 

MnCL + H2SO, = MnSO, + 2 HCl. 

The hydrochloric acid liberated in this reaction is then 
free to act on more of the manganese dioxide, so that all the 
chlorine present in the sodium chloride is set free. 

b. Electrolytic Methods. Chlorine is readily pre- 
pared by the electrolysis of a solution of hydrochloric 
acid in water. The action is very similar to the elec- 
trolysis of water, and can be carried out in the same 
apparatus. It can likewise be obtained by the electro- 
lysis of a- solution of a chloride. In the case of sodium 
chloride for example, the equation representing the re- 
action is NaCl = Na + CI. The sodium however at once 
reacts with the water present in the following way : 
Na + H2O = NaOH + H. The final products of the elec- 
trolysis of a solution of sodium chloride are therefore, 
chlorine which escapes at the anode, hydrogen, which escapes 
at the cathode and sodium hydroxide which remains in solu- 
tion. 

Properties. Chlorine is a greenish yellow gas which 
has a peculiar suffocating odor and a very violent efifect upon 
the throat and lungs. Even when inhaled in small quantities 
it often produces all the symptoms of a hard cold, and in 
larger quantities may have serious and even fatal action. It 
is quite heavy (density=2.45) and can therefore be col- 
lected by downward displacement. It is somewhat sol- 
uble in water, one volume of water under ordinary conditions 



Chlorine Family. 157 

dissolving about three volumes of chlorine gas. The gas is 
readily liquefied, a pressure of six atmospheres serving to 
liquefy it at 0°. It forms a yellowish liquid which solidifies 
to a yellow solid at — 102°. 

Chemical Properties. Chlorine is far more active 
chemically at ordinary temperatures than any of the elements 
we have so far met with, with the exception of fluorine ; in- 
deed it is one of the most active of all elements. 

(i) Action on Metals. A great many metals combine 
directly with chlorine, especially when hot. A strip of cop- 
per foil, heated in a burner flame and then dropped into 
chlorine, burns with incandescence. Sodium burns bril- 
liantly when heated strongly in slightly moist chlorine. Gold 
and silver are quickly tarnished by the gas. 

(2) Action on Non-metals. Chlorine has likewise a 
strong affinity for many of the non-metals. Thus phos- 
phorus burns in a current of the gas while antimony and 
arsenic in the form of a fine powder at once burst into flame 
when dropped into jars of the gas. The products formed 
in all cases where chlorine combines with another element 
are called chlorides. 

(3) Combination zvith Hydrogen. Like oxygen, chlor- 
ine has a strong affinity for hydrogen uniting with it to form 
hydrochloric acid. A jet of hydrogen burning in the air, 
continues to burn when introduced into a jar of chlorine, 
giving a somewhat luminous flame. The two gases when 
mixed together, explode violently when a spark is passed 
through the mixture or when exposed to bright sunlight. 
In the latter case it is the light and not the heat which 
starts the action. 

(4) Action on Substances Containing Hydrogen. Not 
only will chlorine combine directly with hydrogen, but it will 
often abstract hydrogen from chemical compounds which 
contain it. Thus when chlorine is passed into a solution 



158 



Inorganic Chemistry. 



containing hydrosulphuric acid, sulphur is precipitated, and 
hydrochloric acid formed. 

H^S + 2 CI = 2 HCl H- S 

With ammonia the action is similar. 

NH3 + 3 Cl = 3 HCl + N 

The same tendency is very strikingly seen in the action 
of chlorine upon turpentine. The latter substance has the 
composition represented by the formula QoHie- When a 
strip of paper moistened with warm turpentine is placed in 
a jar of chlorine, dense fumes of hydrochloric acid appear, 
and a black deposit of carbon is formed. Even water, which 
is a very stable compound, can be decomposed by chlorine, 
the oxygen being liberated. 

If a long tube of rather large diameter is filled with a 
strong solution of chlorine in water and inverted in a vessel" 

of the same solution as shown in 
Fig. 47, and the apparatus is 
placed in bright sunlight, very 
soon bubbles of a gas will be ob- 
served to rise through the solu- 
tion and collect in the tube. An 
, examination of this gas shows it 
to be oxygen, while the chlorine 
has disappeared from the solution 
and in it's place hydrochloric acid 
is found to be present. 

H20+2Cl=2HCl+0. 
Bleaching Action of Chlorine 
If strips of brightly colored 
cloth, or some highly colored 
flowers are placed in quite dry chlorine, no marked change 
in color is noticed as a rule. If however the cloth and flow- 
ers are first moistened, the color rapidly disappears, that is 
the objects are bleached. Evidently the moisture as well as 
the chlorine is concerned in the action, and a study of .the 



r^ 









^^:=g.-^ 



Fig. 47. 



Chlorine Family, 159 

case shows that the chlorine has combined with the hydrogen 
of the water. The oxygen set free oxidizes the color sub- 
stance converting it into a colorless compound. It is evi-' 
dent from this explanation that chlorine will not bleach all 
colors. Chlorine has also marked germicidal properties^ 
and the free element, as well as compounds from which it 
is easily liberated, are used as disinfectants. 

Nascent State. It will be noticed that oxygen set 
free under these circmnstances is able to do what ordinary 
oxygen cannot do, for both the cloth and the flowers are 
unchanged in the air which contains oxygen. It is generally 
true that elements at the instant of liberation from their com- 
pounds are more active than they are after standing for a 
time in the free state. To express this difference elements 
at the instant of liberation are said to be in the nascent state: 
It is -nascent oxygen which does the bleaching. 

Hydrochloric Acid. The product formed by the 
burning of hydrogen in chlorine is the gas hydrochloric 
acid. This substance is much more easily obtained by treat- 
ing common salt (sodium chloride) with sulphuric acid. 

2 NaCl + H2SO, = Na^SO, + 2 HCl 

Dry salt is placed in a flask furnished with a funnel 
tube and an exit tube, the sulphuric acid is added, and the 
flask gently warmed. Tlie hydrochloric acid gas is rapidly 
given off and can be collected by downward displacement of 
air. The same apparatus can be used as was employed in 
the preparation of chlorine. (Fig. 46.) 

When a solution of salt is treated with sulphuric acid there 
is no very marked action. The hydrochloric acid formed is 
very soluble in water, and so does not escape from the solu- 
tion; hence a state of balance is. soon reached with all four 
of the substances represented in the equation present. Only 
when strong sulphuric acid, in which hydrochloric acid is not 
soluble is poured upon dry salt is the reaction complete. 



i6o Inorganic Chemistry. 

Physical Properties. Hydrochloric acid is a color- 
less gas which has an irritating effect when inhaled, and a 
sour, biting taste, but no marked odor. It is heavier than 
air (density = 1.26) and is very soluble in water. One 
volume of water at 0° and 760 mm. dissolves about 500 
volumes of the gas. On warming such a solution, the 
gas escapes until at the boiling point the solution con- 
tains about 20% by weight of HCl. Further boiling will not 
drive out any more acid, but the solution will distill with 
imchanged concentration. A weaker solution than this will 
lose water on boiling until it has reached the same concentra- 
tion — 20% — and will then distill unchanged. Under high 
pressure the gas can be liquefied, 40 atmospheres being re- 
quired at 10°. It forms a colorless liquid which is not very 
active chemically. Owing to its great attraction for water 
hydrochloric acid fumes strongly in the air. The fumes are 
really minute particles of water which the acid has condensed 
from the air and in which it dissolves. 

Chemical Properties, (i) As the name indicates, 
the substance is an acid, and indeed it is one of the 
strongest of acids ; by this is meant that a large per- 
centage of the acid is dissociated into its ions in solu- 
tions. It acts upon oxides and hydroxides, converting 
them into salts. 

NaOH + HCl = NaCl -f H^O ■ 
CuO + 2 HCl = CUCI2 + HoO 
It acts upon many metals, forming chlorides and liberating 
hydrogen. 

Zn + 2 HCl = ZnClo -f 2 H 
Al -f 3 HCl =:A1C13 -f 3 H 

Unlike nitric and sulphuric acid it cannot oxidize metals, 
so that when it acts at all on them, hydrogen is always given 
ofif. 



Chlorine Family. i6i 

(2) Hydrochloric acid is neither combustible nor a 
supporter of combustion and is therefore not readily decom- 
posed. 

(3) Though hydrochloric acid is incombustible it can 
l3e oxidized under some circumstances, in which case the 
hydrogen combines with oxygen, while the chlorine is set 
free. Thus when hydrochloric acid acts upon manganese 
-dioxide part of the chlorine is set free. 

Mn02 + 4 HCl = MnCl. + 2 B,0 + 2 CI 
Aqua Regia. It has been seen that when nitric acid 
acts as an oxidizing agent it usually decomposes as repre- 
sented in the equation 

2 HNO3 = H,0 + 2 NO + 3 O 
The oxygen so set free is able to oxidize hydrochloric acid : 

6 HCl + 3O = 3 H,0 + 6 CI. 
The complete equation therefore is 

2 HNO3 + 6 HCl = 4 H,0 + 2 NO + 6 CI 
When strong nitric and hydrochloric acids are mixed this re- 
action goes on slowly, chlorine and some other substances 
not represented in the equation being formed. The mixture 
is known as aqua regia ; it acts more powerfully upon metals 
and other substances than either of the acids separately, and 
owes its strength not to acid properties, but to the action 
of the nascent chlorine which it liberates. Consequently 
-A\dien it acts upon metals such as gold, they are converted 
into chlorides, and the reaction can be represented by such 
-equations as 

Au + 3 CI = AuCi, 

Composition of Hydrochloric Acid. When a solu- 
tion of hydrochloric acid is electrolyzed in an apparatus sim- 
iliar to the one in which water was electrolyzed (Fig. 19), 



1 62 Inorganic Chemistry-. 

chlorine collects at one pole and hydrogen at the other. At 
first the chlorine dissolves in the water, but soon the water in 
the one tube becomes saturated with it, and if the stopcocks 
are left open until this is the case, and are then closed, it 
will be seen that the two gases are set free in equal volumes. 
When measured volumes of the two gases are collected over 
a strong salt solution, which does not dissolve either one of 
them, and a spark is passed through the mixture, combina- 
tion takes place, and it is found that one volume of hydro- 
gen combines with one of chlorine. Other experiments show 
that the volume of hydrochloric acid formed is just equal 
to the sum of the volumes of hydrogen and chlorine. One 
volume of hydrogen therefore combines with one volume of 
chlorine to form two volumes of hydrochloric acid gas. 
Since chlorine is 35.18 times as heavy as hydrogen, it fol- 
lows that I part of hydrogen combines with 35.18 parts of 
chlorine to form 36.18 parts of hydrochloric acid. - 

Manufacture of Hydrochloric Acid.. Hydrochloric acid is 
prepared commercially in connection with the manufacture of 
sodium sulphate, which is largely used in the preparation of 
sodium carbonate. Common salt is heated with strong sul- 
phuric acid, and the hydrochloric acid as it escapes in the 
form of gas, is passed into water in which it dissolves, the 
solution being the hydrochloric acid of commerce. When the 
materials are pure, a pure, colorless solution is obtained, the 
strongest solution having a density of 1.2 and containing 40% 
HCl; the commercial acid, often called muriatic acid, is usually 
colored yellow with impurities. 

Compounds of Chlorine with Oxygen and Hydrogen, 

Chlorine combines with oxygen and hydrogen to form no 
less than four different substances, all of which have the 
properties of acids. They are all quite unstable, and most 
of them cannot be prepared in pure form. Their salts can 
easily be made however, and some of them will be met with 



Chlorine Family. 163 

in the study of the metals. The names and formulas of these 
acids are as follows : 

HCIO, hypochlorous acid 

HCIO2, chlorous acid. 

HCIO3, chloric acid 

HCIO4, perchloric acid 

Oxides of Chlorine. Two oxides are known, having 
the formulas CI2O and CIO2. They decompose very easily, 
and are good oxidizing agents. 

BROMINE. 

Historical. Bromine was discovered in 1826 by the 
French chemist Ballard, who isolated it from sea salt. He 
named it bromine (stench), because of its unbearable fumes. 

Occurrence. Bromine occurs almost entirely in the 
form of bromides, especially as sodium bromide and mag- 
nesium bromide which are found in many salt springs and 
salt deposits. The Stassfurt deposits in Germany, and the 
salt waters of Michigan and Ohio are especially rich in 
bromides. 

' Preparation of Bromine. 

* I. Laboratory Method. As in the case of chlorine, bro- 
mine can be prepared by the action of hydrobromic acid^ 
HBr, on manganese dioxide. Since the hydrobromic acid 
is not an article of commerce, a mixture of sulphuric acid 
and a bromide is commonly substituted for it. The materials 
are placed in a retort like the one used in the preparation of 
nitric acid (Fig. 33). On heating the bromine distills over 
and is collected in the cold receiver. 

2. Commercial Method. Bromine is prepared com- 
mercially from the waters from salt wells, those being taken 
which are rich in bromides. On passing a current of elec- 



164 Inorganic Chemistry. ^ 

tricity through such waters, the bromine is first liberated. 
-Any chlorine liberated however will assist in the reaction, 
.since free chlorine decomposes bromides as shown in the 
^equation 

NaBr + CI = NaCl + Br 

Properties. Bromine is a dark red liquid, about three 
times as heavy as water. Its vapor has a very offensive odor, 
and is most irritating to the eyes and throat. The liquid 
boils at 59° and solidifies at — 7° ; but at ordinary temper- 
atures it evaporates rapidly, forming a reddish brown gas, 
very similar to nitrogen peroxide in appearance. Bromine 
:is somewhat soluble in water, 100 volumes of water under 
uDrdinary conditions dissolving i volume of the liquid. 

In chemical action bromine is very similar to chlorine. 
It combines directly with many of the same elements with 
which chlorine unites, but with less energy. It combines 
with hydrogen and takes away the latter element from its 
compounds, but not so readily as does chlorine. Its bleach- 
ing properties are also less marked. 

Hydrobromic Acid. HBr. When sulphuric acid acts 
upon a bromide, hydrobromic acid is set free. 

2 NaBr + H^SO, = Na^SO^ + 2 HBr 
At the same time some bromine is set free, as can be seen 
from the red fumes which appear, and from the odor. The 
explanation of this is found in the fact that hydrobromic acid 
is much less stable than hydrochloric acid, and is therefore 
more easily oxidized. Sulphuric acid is a good oxidizing 
agent. So the strong sulphuric acid oxidizes a part of the 
hydrobromic acid, forming bromine from it. 
'H2SO4 + 2 HBr = 2 H2O + SO2 + 2 Br 

A convenient way to make pure hydrobromic acid is by 
the action of bromine upon moist red phosphorus. This can 



Chlorine Family. 



165 




Fig. 48. 



be done with the apparatus shown in Fig. 48. Bro- 
mine is put into the dropping funnel A, and red phos- 
phorous, together with enough water to cover it, intO' 
the flask B. By means of the stopcock, the bromine 

is allowed to flow drop by 
drop into the flask, the re- 
action taking place with- 
out the application of heat. 
The equations are: 

P + 3Br = PBr3 
PBrs + sH^O^ 
P(OH)3 + 3HBr 
The U-tube C contains 
glass beads which have been moistened with water and rubbed 
in red phosphorus. Any bromine escaping action in the flask 
acts upon the phosphorus in the U-tube. The hydrobromic 
acid is collected in the same way as hydrochloric acid. 

Properties. Hydrobromic acid very strikingly re- 
sembles hydrochloric acid in physical and chemical proper- 
ties. It is a colorless, strongly fuming gas, heavier than 
hydrochloric acid, and like it, very soluble in water. One 
volume of water under standard conditions dissolves 610 
volumes of the gas. Chemically, the chief point in which it 
differs from hydrochloric acid is in the fact that it is much 
more easily oxidized, so that bromine is more easily set free 
from it than chlorine is from hydrochloric acid. 

The bromides are very similar to the chlorides in their 
properties. Chlorine acts upon both bromides and free hy- 
drobromic acid, liberating bromine from them. 
KBr + CI = KCl + Br 
HBr + CI = HCl + Br 

Bromine finds many uses in the manufacture of organic 
drugs and dye stuffs. Some of the bromides are used in 
photography and as drugs. 

Oxygen Compounds. No oxides of bromine are 
surely known, and bromine does not form so many oxygen 



1 66 Inorganic Chemistry. 

acids as chlorine does. Salts of hypobromous acid, HBrO, 
and bromic acid, HBrOg, are known. 

IODINE. 

Historical. Iodine was discovered by Courtois 
(1812) in the ashes of certain sea plants. Its presence was 
revealed by its beautful violet vapor, and this suggested the 
name iodine (from the Greek for violet appearance). 

Occurrence. In the combined state iodine occurs in 
very small quantity in sea water, from which it is absorbed 
by certain sea plants, so that it is found in their ashes. It 
occurs along with bromine in salt springs and beds, and 
is also found in small amounts in Chili saltpetre. 

Preparation. Iodine can readily be prepared in the 
laboratory from an iodide by the same method as is used 
in preparing bromine, only that in place of sodium bromide, 
sodium iodide is substituted. It can also be made by passing 
chlorine into a solution of an iodide. 

Commercially, iodine was formerly prepared from sea- 
weed (kelp) but is now obtained almost entirely from the 
mother liquors from which Chili saltpetre has been crystal- 
lized. The iodine is present in the liquors in the form of 
sodium iodate, and the chemical reaction by which it is pre- 
pared is a complicated one. It depends on the fact that sul- 
phurous acid acts upon iodic acid setting iodine free. 

2 HIO3 + 5 H,S03 =: 5 H,SO, + H,0 + 2 I 

Physical Properties. Iodine is a purplish-black, 
shining, heavy solid which crystallizes in brilliant plates. 
Even at ordinary temperatures it gives off a beautiful violet 
vapor, which increases in amount as the iodine is heated. 
It melts at 107° and boils at 175°. It is shghtly soluble in 
water, but readily dissolves in alcohol, forming a brown so- 
lution (tincture of iodine), and in carbon disulphide, form- 
ing a violet solution. When even minute traces of it are 



Chlorine Family. 167 

added to thin starch paste, a very intense blue color de- 
velopes and this reaction forms a delicate test for iodine. 
The element has a strong, unpleasant odor, though by no 
means so irritating as that of chlorine and bromine. 

Chemical Properties. Chemically, it is quite similar 
to chlorine and bromine but is still less active than bromine. 
It combines directly with many elements, and at elevated 
temperatures with hydrogen ; but the latter reaction is in- 
complete and the compound formed is quite easily decom- 
posed. Both chlorine and bromine displace it from its salts. 
KI + Br = KBr + I 
KI + CI = KCl + I 

Hydriodic Acid. Hydriodic acid cannot be prepared 
in pure form by treating an iodide with sulphuric acid, since 
tinder these conditions iodine and not hydriodic acid is set 
free. The explanation of this is the same as in the case of 
b)romine. Hydriodic acid is a strong reducing agent, and 
when set free, acts upon the sulphuric acid. 

2 Nal + H,SO, = Na.SO, + 2 HI 
2 HI + H,s6, = 2 H^O + SO, + 2 I 
It is most easily prepared in the way described for the prepa- 
ration of hydrobromic acid. Red phosphorus and iodine 
are rubbed together and placed in a small flask. The rriix- 
ture is warmed gently, and water is slowly dropped upon the 
mixture. 

P + 3 I = PI3 
Pl3 + 3H,0 = 3HI + P(OH)3 

Physical Properties. Hydriodic acid is a strongly 
fuming colorless gas which might be easily mistaken for 
either hydrochloric or hydrobromic acids. It is very solu- 
ble in water, i volume of water under standard conditions 
dissolving about 460 volumes of the gas. It is easily con- 
densed to a colorless liquid by a pressure of four atmos- 
pheres at 0°. 



1 68 Inorganic Chemistry. 

Chemical Properties. A solution of hydriodic acid 
has strong acid properties, acting- very much like the cor- 
responding solutions of hydrochloric and hydrobromic acids. 
It is however much more unstable. On standing exposed 
to the air, the solution turns brown from the separation 
of iodine ; 

2 HI + O = H,0 + 2 I 
Owing to the small affinity between the iodine and the hy- 
drogen, the compound gives up its hydrogen very easily, 
and is therefore a very strong reducing agent. Of the salts 
of this acid, potassium iodide is the most familiar. 

• Iodine has a much greater affinity for oxygen than has 
either chlorine or bromine. It forms a stable oxide I2O5 
which is made by boiling iodine with nitric acid. Salts of 
iodic acid, HIO3 and periodic acid, HIO4, are easily pre- 
pared, and the free acids are much more stable than those 
of the other members of this family. 

Uses of Iodine. Iodine is extensively used in. medi- 
cine, especially in the form of a tincture, and it is also 
much used in the preparation of dyes and organic drugs. 
Iodoform, a substance with the composition CHI3, is much 
used in surgery as an antiseptic. 



• CHAPTER XV. 

CARBON AND SOME OF ITS SIMPLER COMPOUNDS. 

Occurrence. Carbon is found in nature in the un- 
combined state in several forms. The diamond is prac- 
tically pure carbon, while graphite and coal are largely car- 
bon, but contain small amounts of other substances mixed 
with it. Its natural compounds are exceedingly numerous, 
and occur as gases, liquids and solids. Carbon dioxide 
is its most familiar gaseous compound. Natural gas and 
petroleum are largely compounds of carbon with hydrogen. 
The carbonates, especially calcium carbonate, constitute great 
strata of rocks, and are found in almost every locality. All 
living organisms, both plant and animals, contain a large 
percentage of this element, and the number of its compounds 
which go to make up all the vast variety of animate nature 
is almost limitless. 

Allotropic Forms. Carbon occurs in three distinct 
forms, two of which are crystalline and one amorphous. 

I. Crystalline Forms. 

a. Diamond. Diamonds are found in considerable 
quantities in several localities, especially in the East Indies, 
South Africa and Brazil. The crystals belong to the 
regular system, but the natural stones do not show this 
very clearly. When found they are usually covered with 
a rough coating which is removed in the process of cutting. 

The pure diamond is perfectly transparent and colorless, 
but many are tinted a variety of colors by traces of metallic 
substances. In some instances the color adds to the value; 
thus the famous Hope diamond is a beautiful blue. Usually 
the colorless ones are the most highly prized. Light passing 
through a diamond is very much refracted, and to this fact 
the stone owes its brilliancy and sparkle. 

(169) 



170 Inorganic Chemistry. 

The density of the diamond is 3.5 and though brittle, 
it is one of the hardest of substances. Black diamonds, as well 
as broken and imperfect stones, which are valueless as gems 
are used for grinding hard substances. Few chemical re- 
agents have any action on the diamond ; but when heated 
strongly in oxygen or the air, it blackens and burns forming 
carbon dioxide. 

Lavoisier first showed that carbon dioxide is formed by 
the combustion of the diamond ; and Sir Humphrey Davy in 
1814 showed that this is the only product of combustion, 
and that the diamond is therefore pure carbon. 

Many attempts have been made to produce diamonds ar- 
tificially, but for a long time these always ended in failure, 
graphite and not diamonds being the product obtained. In 
liis extended study of chemistry at high temperatures, the 
French chemist Moissan finally succeeded (1893) in making 
some small ones. He accomplished this by dissolving carbon 
in boiling iron and suddenly plunging the crucible containing 
the iron into water. Under these conditions the carbon was 
crystallized in the iron in the form of the diam.ond. 

b. Graphite. Graphite is found in large quantities es- 
pecially in Ceylon, Siberia and in some localities of the 
United States and Canada. It is a shining black substance, 
very soft and greasy to the touch. Its density is about 2.15. 
It varies somewhat in properties according to the locality in 
which it is found, and is more easily attacked by reagents 
than is the diamond. It is manufactured by heating carbon 
with a small amount of iron (3%) in an electric furnace. 
It is used in the manufacture of lead pencils and crucibles, 
as a lubricant where oil cannot be used, and as a protective 
covering for iron in the form of a polish, such as stove polish, 
or as a paint. 

2. Amorphous Carbon. There are many varieties of 
amorphous carbon known, and these substances are of the 
greatest importance, owing to their many uses in the arts 



Carbon and Some of its Simpler Compounds. lyi 

and industries. As they occur in nature, or are made arti- 
iicialiy, they are nearly all impure carbon, the impurity de- 
pending on the particular substance in question. 

a. Pure amorphous carbon is best prepared by charring 
sugar. This is a substance consisting of carbon, hydrogen 
and oxygen, the latter two elements being present in the 
ratio of one oxygen atom to two of hydrogen. When sugar 
is strongly heated, the oxygen and hydrogen are driven off 
in the form of water, and pure carbon is left behind. Pre- 
pared in this way it is a soft, lustrous, very bulky, black 
powder. 

b. Coals of various kinds were probably formed from 
vast accumulations of vegetable matter in former ages, 
which became covered over with earthy material and pro- 
tected from rapid decay. Under various natural agencies 
the organic matter was slowly changed into coal. In an- 
thracite, these changes have gone the farthest, and this va- 
riety of coal is nearly pure carbon. Soft, or bituminous, 
coals contain considerable organic matter besides carbon 
and mineral substances. When distilled, (heated strongly out 
of contact with air), the organic matter is driven off in the 
form of gases and vapors, and only the mineral matter and 
carbon remain behind. The gaseous product is chiefly il- 
luminating gas, and the solid residue is coke. Some of the 
coke is found as a dense cake on the sides and roof of the 
retort. This is called retort carbon and is quite pure. 

c. Charcoal is prepared from wood in the same way 
that coke is made from coal. When the process is carried 
on in retorts, the products expelled by the heat are saved. 
Among these are many valuable substances such as wood al- 
cohol and acetic acid. Where timber is abundant, the pro- 
cess is carried out in a wasteful way, by merely covering piles 
of wood with sod, and setting the wood on fire. Some wood 
burns, and the heat from this decomposes the wood not 



172 Inorganic Chemistry. 

burned, forming charcoal from it. The charcoal of course 
contains the mineral part of the wood from which it is 
formed. 

d. Bone-black or anirhal charcoal is made by charring 
bones and animal refuse. The organic part of the materials 
is thus decomposed, and carbon is left in very finely divided 
state, scattered through the mineral part, which consists 
largely of calcium phosphate. For some uses this mineral 
part is removed by treatment with hydrochloric acid, and 
prolonged washing. 

e. Lampblack and soot are products of imperfect com- 
bustion of oil and coal, and are deposited from a smoky 
flame on a cold surface. The carbon in this form is very 
finely divided, and usually contains various oily materials. 

Properties. Carbon in any of its forms is charac- 
terized by its great stability towards heat. Only in the in- 
tense heat of the electric arc does it volatilize, passing di- 
rectly from the solid state into a vapor. It is not acted on at 
ordinary temperatures by very many reagents, but at a higher 
temperature combines with many" of the elements. Its com- 
pounds with the metals are called carbides. 

Uses of Carbon. The chief use of amorphous car- 
bon is for fuel to furnish heat and power for all the uses of 
civilization. Enormous quantities of the purer coals, coke 
and charcoal are used as a reducing agent in the manufac- 
ture of the various metals, especially in the metallurgy of 
iron. Thus iron is found in nature chiefly in the form of an 
oxide. When heated with carbon, this oxide is reduced to 
metallic iron. 

This use can be readily illustrated by heating in a hard 
glass test tube an intimate mixture of 2 or 3 grams of copper 
o^ide with an equal bulk of powdered charcoal. The mixture 
turns reddish and if the evolved gases are passed through clear 
lime water, the presence of carbon dioxide can be demonstrated. 



Carbon and Some of its Simpler Compounds. 173 

Retort carbon and coke are used to make electric light 
carbons and battery plates, while lampblack is used for 
indelible inks, printer's ink and black varnishes. Boneblack 
and charcoal have the property of absorbing large volumes 
•of certain gases, as well as smaller amounts of organic 
matter; hence they are used in filters to remove noxious 
g-ases and objectionable colors and odors from water. Bone- 
black is used extensively in the sugar refineries to remove 
coloring matter from the impure sugars. 

Fuels. A variety of substances are used as fuels, the 
most important of them being coal, wood and gas. These 
are composed of carbon, hydrogen and, in the case of the 
first two, oxygen. When the fuel burns in a free supply of 
air, the carbon is converted into carbon dioxide, and the 
hydrogen into water. In the case of gas, which often con- 
tains a large percentage of hydrogen, the amount of water 
vapor formed is very great, and rooms heated by gas fires 
are often very damp because of faulty ventilation for the 
products of combustion of the gas. 

Chemistry of Carbon Compounds. Carbon is re- 
markable for the very large number of compounds which it 
forms with the other elements, especially with oxygen and 
hydrogen. Compounds containing carbon are more numer- 
ous than all others put together, and the chemistry of these 
substances presents peculiarities not met with in the study 
of other substances. For these reasons the systematic study 
of carbon compounds or of organic chemistry as it is usually 
called, must be deferred until the student has gained some 
knowledge of the chemistry of other elements. An ac- 
quaintance with a few of the most familiar carbon com- 
pounds is, however, essential for the understanding of the 
general principles of chemistry. 

Compounds of Carbon with Hydrogen. Carbon 
forms a very large number of compounds with hydrogen, 



174 Inorganic Chemistry. 

and when no other element is present, such substances are 
called hydrocarbons. Petroleum and natural gas are es- 
sentially mixtures of a great variety of hydrocarbons. Many 
are found in living plants, and the decay of organic mat- 
ter in the absence of air always produces them. 

Marsh Gas — Methane. Marsh gas, CH4 is one of 
the most important of these hydrocarbons. As its name sug- 
gests it is formed in marshes by the decay of vegetable mat- 
ter under water, and bubbles of it are often seen to rise 
when the dead leaves on the bottom of pools are stirred. It 
also collects in mines, and is called fire damp by the miners, 
because of its great inflammability — damp being an old name 
for a gas. It is formed when organic matter, such as coal 
or wood, is heated in closed. vessels, and is therefore a prin- 
cipal constiuent of coal gas. 

Methane is prepared in the laboratory by heating a mix- 
ture of sodium or calcium acetate and soda-lime. Equal weights 
of fused sodium acetate and soda-lime are thoroughly mixed,, 
and placed in a good-sized, hard-glass test tube, fitted with a 
one-holed stopper and delivery tube. The mixture is grad- 
ually heated, and when the air has been displaced from the 
tube, the gas is collected in bottles by displacement of water. 
Soda lime is a mixture of sodium and calcium hydroxides. Re- 
garding it as sodium hydroxide alone, the equation is 
NaC2H302+NaOH=Na2C03+CH4. 

Properties. Methane is a colorless, odorless gas, 
whose density is 0.55. It is difficult to liquefy, boiling at 
— 155° under standard pressure, and is almost insoluble fn 
water. It burns with a pale blue flame, liberating much 
heat, and mixed with oxygen it forms a very explosive 
mixture. 

Davy's Safety Lamp. In 181 5 Sir Humphrey Davy in- 
vented a lamp for the use of miners to prevent the dreadful 
mine explosions then common, due to methane mixed with 
the oxygen of the air. The invention consisted in surround- 
ing the common miner's lamp with a mantle of wire gauze. 



Carbon and Sojne of its Simpler Compounds. 175 



It has been seen that two gases will not combine until raised 
to their kindling temperature, and if, while combining, they 
are cooled below this point, the combination ceases. A flame 
will not pass through a wire gauze because the metal, being 
a good conductor of heat, takes away so much heat from the 
flame that the gases are cooled below the kindling tempera- 
ture. When a lamp so protected is brought into an explosive 
mixture, the gases inside the wnre mantle burn in a series 
of little explosions, giving warning to the miner that the air 
is unsafe. 

Acetylene, C^Ho,, is a colorless gas usually possessed 
of a disagreeable odor due to impurities. It is formed in 
many processes, and is one of the most important constitu- 
ents of illuminating gas. It is now made in large quantities 
from calcium carbide, CaCo. This substance is formed when 
coal and lime are heated together in an electric furnace. 
When treated with water, the carbide is decomposed yield- 
ing acetylene. 

CaCo + 2 H2O = QH2 + Ca(OH)2. 
The gas can be readily liquefied, but it is not used in this 
way, because it has been found that the liquid sometimes 
explodes especially when exposed to a sudden shock. Under 
ordinary conditions it burns with a very smoky 
flame; in burners constructed so as to secure a 
large amount of oxygen, it burns with a very 
brilliant white light, and hence is used as an 
illuminant. The gas can be readily prepared in 
a generator such as is shown in Fig. 49. The 
inner tube contains fragments of calcium car- 
bide, the outer, water. The gas can be lighted 
as it issues from the burner. 

Other Hydrocarbons. At first thought 
the law of multiple proportion would seem to 
indicate that there must be a limit to the num- 
ber of compounds which two elements can form 




Fig. 



176 Inorganic Chemistry. 

with each other, and our experience so far has shown that 
in general the number is a small one. Yet hundreds of hydro- 
carbons are known. The ability of carbon to form such a 
surprising number of compounds with hydrogen appears to 
be due to the fact that carbon atoms have the peculiar 
power of combining with each other to form long chains 
and this makes possible a great variety of compounds. Thus 
we have the hydrocarbons 

Ethane H3C — CH3 

Propane H3C — CH^ — CHo 

Butane H3C — CHo — CH^ — CH3 

Oxides of Carbon. Carbon forms two oxides, 
viz., carbon dioxide, CO2, and carbon monoxide, CO. 

I. Carbon Dioxide COg. Occurrence. Carbon diox- 
ide is present in the air to the extent of about 4 parts in 
10,000, and this apparently small amount is of fundamental 
importance in nature. It escapes from the earth in some 
localities in great quantities, and many spring waters carry 
large amounts of it in solution. When these highly charged 
spring waters reach the surface of the earth, and the pressure 
on them is removed, the carbon dioxide escapes with effer- 
vescence. It is a product of the oxidation of all organic 
matter and is therefore formed in fires, as well is in the 
process of decay. It is thrown off from the lungs of all 
animals in respiration. It is a product of many fermentation 
processes such as vinegar making and brewing. Combined 
with metallic oxides, it forms vast deposits of carbonates 
in nature. 

Preparation. In the laboratory, carbon dioxide is 
always prepared by the action of an acid upon a carbonate, 
usually calcium carbonate, the apparatus shown in Fig 42 
serving the purpose very well. This reaction might be ex- 
pected to produce carbonic acid, thus : 

CaC03 + 2 HCl = CaCU + H^COg 



Carbon and Some of its Simpler Compounds. 177 

Carbonic acid is very unstable however, and decomposes into 
its anhvdride, CO2, and water. 

H2CO3 = H2O + COo. 
The complete reaction is therefore represented by the equa- 
tion : 

CaC03 + 2 HCl = CaCl^ + CO, + H^O. 

The reaction is therefore quite similar to the one in which 
sulphur dioxide is prepared by the action of an acid on a 
sulphite : 

Na.SOa + 2 HCl = 2 NaCl + SO, + H,0. 
Physical Properties. Carbon dioxide is a colorless, 
practically odorless gas whose density is 1.5. Its weight can 
be inferred from the fact that it can be siphoned or poured 
hke water from one vessel downward into another vessel. It 
is somewhat soluble in water. At 15° and under ordinary 
pressure, it dissolves in its own volume of water, and imparts 
a somewhat biting,- pungent taste to it. It is easily con- 
densed and is prepared now commercially in this form, by 
pumping the gas into steel cylinders (see Fig. 6), which are 
kept cold during the process. When the liquid is permitted 
to escape into the air, part of it instantly evaporates and in 
so doing, absorbs so much heat that another portion is solidi- 
fied, the solid form strikingly resembling snow in appear- 
ance. This snow is very cold, and mercur}^ can easily be 
frozen with it. 

Such cylinders of carbon dioxide are inexpensive, and 
should be available in every school. To demonstrate the prop- 
erties of solid carbon dioxide, the cylinder should be placed 
across the table and supported in such a way that the stopcock 
end is several inches lower than the bottom of the cylinder. A 
loose bag is made by holding the corners of a handkerchief 
around the neck of the stopcock, and the cock is then turned on 
so that the gas rushes out in large quantities. In a few minutes 
a considerable quantity of the snow collects in the handker- 
chief. To freeze mercury, press a piece of filter paper into a 



178 Inorganic Chemistry. 

small evaporating dish, and pour the mercury upon it. Coif 
a flat spiral upon the end of a wire, and dip the spiral inta 
the mercury. Place a quantity of solid carbon dioxide upon 
the mercury and pour 10-15 c.c. of ether over it. In a minute 
or two the mercury will solidify, and can be removed from, 
the dish by the wire serving as a handle. The Alter paper 
is to prevent the mercury sticking to the dish; the ether dis- 
solves the solid carbon dioxide and promotes its rapid con- 
version into gas. 

In the liquid form carbon dioxide is sold for charging" 
water with the gas — such highly charged water being used 
in soda water and similar drinks. 

Chemical Properties. Carbon dioxide is incombus- 
tible, since it is like water a product of combustion. It does 
not support combustion, as does nitrogen peroxide because 
the oxygen in it is held in very firm chemical union with 
the carbon. Very strong reducing agents such as highly 
heated carbon, can take away half of its oxygen : 
CO, + C = 2 CO. 

Carbon dioxide and water are therefore both of them 
good tire extinguishers. Chemical tire extinguishers are simply 
devices for rapidly generating a large amount of carbon dioxide. 
The gas is directed upon the burning substance and smothers 
the fire just as water would. 

It is not necessary that all the oxygen should be kept 
away from the fire to smother it. A burning candle is ex- 
tinguished in air which has been mixed with only 2.5% carbon 
dioxide, and this fact adds very much to the effectiveness of 
fire extinguishers which make use of carbon dioxide. 

When passed into clear lime water, carbon dioxide 
forms an insoluble substance, calcium carbonate, and this is 
the easiest way of testing for the gas. 

Carbonic Acid and the Carbonates. Like most of 
the oxides of the non-metallic elements, carbon dioxide is 
afi acid anhydride. It combines with water to form an acid 
of the formula H2CO3, called carbonic acid. 
H2O -f CO, = H2CO3. 



Carbon and Some of its Simpler Compounds. 179 

The acid is, however, very unstable, and cannot be iso- 
lated. Only a very small amount of it is actually formed 
when carbon dioxide is passed into water, as is evident from 
the small solubility of the gas. If, however, a base is present 
in water, salts of carbonic acid are formed, and these are 
quite stable. 

2 NaOH + H.O + CO, = NaXOg + 2 H^O. 

This conduct is best explained by the law of mass action 
already referred to. The equation H20+C02=H2C03, is a re- 
versible equation, and the extent to which the reaction pro- 
gresses depends upon the relative amounts — that is, upon the 
active masses, of each of the three factors in it. Equilibrium 
is ordinarily reached when very little H2CC33 is formed. . If 
the quantity of CO2 is increased by pressure, more of the 
acid is formed; and when the pressure is" removed the re- 
action is reversed, and the CO2 is given off. If a base is 
present in the water to take up the H2CO3 as fast as • it is 
formed, that is to keep the active mass of the acid very smalh 
all of the CO2 is converted into H2CO3, and thence into a 
carbonate. When an acid which contains a gaseous anhy- 
dride, such as H2CO3, H2SO3, HNO2, is set free, it is apt to 
decompose, and if its anhydride is insoluble or nearly so, the 
decomposition is complete, because one of the factors in the 
reaction is constantly escaping, and this reduces its active 
mass below the point of equilibrium. 

The carbonates form a very important class of salts. 
They are found in large quantities in nature, and are often 
used in chemical processes. Only the carbonates of sodium, 
potassium and ammonium are soluble, and these can be made 
by the action of carbon dioxide on solutions of the bases as 
just explained. 

The insoluble carbonates are formed as precipitates, 
when a soluble metallic compound is treated with a solution 
of a soluble carbonate. Thus the insoluble calcium car- 
bonate can be made by bringing together solutions of cal- 
cium chloride and sodium carbonate: 

CaCl^ + Na^COg = CaCOg -f 2 NaCl 



i8o Inorganic Chemistry. 

Most of the carbonates are decomposed by heat, yielding an 
oxide of the metal and carbon dioxide. Lime (calcium ox- 
ide) is made by strongly heating calcium carbonate: 
CaCOg = CaO + CO^. 
Acid Carbonates. Like all acids containing two acid 
hydrogen atoms, carbonic acid can form both normal and 
acid salts. The acid carbonates are made by treating a nor- 
mal carbonate with an excess of carbonic acid. They are 
all very unstable, and decompose even when their solutions 
are heated. 

If carbon dioxide is passed into clear lime water, calcium 
carbonate is at first precipitated: Ca(OH)2+C02=CaC03+H20. 
The precipitate soon dissolves, because the excess of car- 
bonic acid forms calcium acid carbonate: CaC03+H2C03= 
Ca(HC03)2. If now the solution is heated, the acid carbonate 
will be decomposed, and calcium' carbonate once more pre- 
cipitated. Ca(HC03)2=CaC03+H2C03. 

Carbon Dioxide and Plant Life. Although large 
amounts of carbon dioxide are constantly being formed in 
nature, the amount in the air remains almost constant. The 
reason for this is that plant life is dependent upon carbon 
dioxide for food, and that growing plants absorb it from 
the air. Although it is a very stable substance, it is decom- 
posed in the plant under the influence of the sunlight. Water 
also undergoes decomposition at the same time. Part of the 
oxygen of these two substances is restored to the air from the 
leaves of the plants, while the carbon and hydrogen and a 
part of the oxygen are biiilt up into the complex structure 
of the plant. The plants are therefore laboratories where 
■carbon dioxide and water are made over into more complex 
substances, while the oxygen is in part returned to the air. 

Since much heat is liberated when wood is burned, that 
Is when the process of its formation is reversed, the oxygen 
iDcing taken again from the air and the wood being con- 



Carbon and Some of its Simpler Compounds. i8i 

verted into carbon dioxide and water, it is evident that 
much energy is absorbed when the carbon dioxide and water 
are decomposed and built up into the woody structure. This 
energy is suppHed by the sun acting on the growing plant. 
Thus the plants serve as reservoirs to store up for a time 
the energy of the sun in the form of complex chemical sub- 
stances, and this energy is again liberated when the plant is 
burned. 

2. Carbon Monoxide. Carbon monoxide can be made 
in two general ways : ( i ) by partial oxidation of carbon, 
(2) by partial reduction of carbon dioxide. If a slow cur- 
rent of air is conducted over highly heated carbon, the mon- 
oxide is formed : C + O ^ CO. It is therefore often formed 
in stoves when the air draught is insufficient. Water gas, 
w^hich contains large amounts of carbon monoxide, is made 
by partially oxidizing carbon with steam : 

C + H,0 = CO + H, 

When COo is conducted over highly heated carbon, it is re- 
duced to CO by the excess of carbon : COo -\- C ^= 2 CO. 
When coal is burhing in a stove or grate, carbon dioxide is 
at first formed in the free supply of air ; as the hot gas 
rises through the glowing coal, it is reduced to carbon mon- 
oxide. When the carbon monoxide reaches the free air 
above the coal, it takes up oxygen to form carbon dioxide, 
burning with the blue flame so familiar above a bed of coals, 
especially in the case of hard coal. In the laboratory carbon 
monoxide is usually prepared by the action of sulphuric acid 
upon oxalic acid. The latter substance has the formula 
C2H2O4. The sulphuric acid owing to its affinity for water 
causes it to decompose as represented in the equation 

C,U,0, + (H^SOJ = (H.SOJ + H^O + CO^ + CO 

Properties. Carbon monoxide is a light, colorless, al- 
most odorless gas, very difficult to liquefy. Chemically it is 



1 82 Inorganic Chemistry. 

very active, combining directly with a great many sub- 
stances. It has a great affinity for oxygen, and is therefore 
combustible, and a good reducing agent. Thus if carbon 
monoxide is passed over hot copper oxide, the copper is re- 
duced to the metallic state : CuO + CO = Cu + COg. 
When inhaled, it combines with the red coloring matter of 
the blood, and even a very small amount of the gas will cause 
•death. 

Carbon Bisulphide. Just as carbon combines with 
oxygen to form carbon dioxide, so it combines with .sul- 
phur to form carbon disulphide, CS2. This compound has 
already been described in the chapter on sulphur. 

Hydrocyanic Acid — Cyanides. Carbon unites with 
nitrogen and hydrogen to form the acid HCN, called 
hydrocyanic acid. It is a weak volatile acid, and is there- 
fore easily prepared by treating its salts with sulphuric 
acid. 

KCN + H.SO^ = KHSO, + HCN. 
It is most familiar as a gas, though it condenses to a color- 
less liquid, boiling at 26°. It has a peculiar odor suggesting 
bitter almond, and is extremely poisonous. A single drop may 
cause death. It dissolves readily in water, its solution being 
commonly called prussic acid. Its salts are called cyanides, 
the cyanides of sodium and potassium being the best known. 
These are white solids and extremely poisonous. 



CHAPTER XVI. 

FLAMES — ILLUMINANTS. 

It has been seen that \vhen two substances unite chemi- 
cally, with the production of light and heat, the act of union 
is called combustion. When one of the substances under- 
going combustion remains solid at the temperature occa- 
sioned by the combustion, light may be given off, but there 
is no flame. Thus, iron wire burning in oxygen throws off 
a shower of sparks, and is brilliantly incandescent, but no 
flame is seen. When however both of the substances are 
^ases or vapors at the temperature reached in the combus- 
tion, the act of union is accompanied by a flame. 

]\Iany substances which are liquids or solids at ordinary- 
temperatures burn with a flame because the heat of combus- 
tion vaporizes them slowly, and the flame is due to the union 
of this vapor with the gas supporting the combustion. Thus 
the oil in a lamp is drawn up through the wick by capillary 
attraction and is vaporized by the heat of the flame. Again 
in the burning of a candle the heat of combustion is sufficient 
to slowly melt the solid wax at the end of the candle. The 
resulting liquid is drawn up in the wick by capillary attrac- 
tion and is vaporized, the flame being due to the union of this 
vapor with oxygen in the air. Solid substances like wood burn 
with a flame because the heat of combustion expels volatile 
combustible inatter from the wood. If the wood is first heated 
strongly in the absence of the air, the volatile matter is all 
•expelled and the resulting substance, charcoal, simply glows 
when it burns, there being no flame. 

Supporter of Combustion. That gas w^hich sur- 
rounds the flame and constitutes the atmosphere in which 
the combustion occurs is said to support the combustion. 
The other gas which issues into this atmosphere is said to 
be the combustible gas. Thus in the ordinary combustion 

(183) 



1 84 



Inorganic Chemistry. 



of coal gas in the air, the coal gas is said to be combustible 
while the air is the supporter of combustion. These terms 
are entirely relative however for a jet of air issuing into an 
atmosphere of coal gas will burn when ignited, the coal gas 
supporting the combustion. Ordinarily when we say that a 
gas is combustible, we mean that it is combustible in an at- 
mosphere of air. 

That the terms "combustible" and "supporter of combus- 
tion" are merely relative may be shown in the following way. 
A lamp chimney is fitted with a cork and 
glass tubes as shown in Fig. 50. The tube C 
should have a diameter of from 12 to 15 mm. 
A thin sheet of asbestos in which is cut a 
circular opening about 2 cm. in diameter, is 
placed over the top of the chimney. The 
opening in the asbestos is closed with the 
palm of the hand and gas is admitted to the 
chimney through the tube B. The air in the 

^ . chimney is soon expelled through the tube 

^*^ C and the gas itself is then lighted at the 

end of this tube. The hand is now removed 
from the opening in the asbestos, when the 
flame at the end of the tube at once rises 
and appears at the end of the tube within 
the chimney as shown in the figure. The coal gas now escapes 
from the opening in the asbestos and may be lighted. The 
flame at the top of the chimney is due to the combustion of 
coal gas in air while the flame within the chimney is due to 
the combustion of air in coal gas, the air being drawn up 
through the tube by the escaping coal gas. 

Appearance of Flames. The flame caused by the 
union of hydrogen and oxygen is almost colorless and in- 
visible. Chlorine and hydrogen combine with a .pale vio- 
let flame ; carbon monoxide burns in oxygen with a blue 
flame, ammonia with a deep yellow flame. The color and 
.appearance of flames are therefore often quite characteristic 
of the particular combustion which occasions them. 






Fig. 50. 



Flames — Illiuninants. 185 

Structure of Flames. When the gas undergoing 

combustion issues from a round opening into an atmosphere 

of the gas supporting the combustion, as is the case with the 

burning Bunsen burner shown in Fig. 51, the 

A---B flame is in general conical in outline. It con- 
sists of several distinct cones, one within the 
other, the boundary between them being marked 
by differences of color or luminosity. In the 
simplest flame, of which hydrogen burnmg in 
oxygen is a good example, these cones are two 
in number; an inner one, formed by unburned 
gas, and an outer one, usually more or less 
luminous, consisting of the combining gases. 
This outer one is in turn surrounded by a third 
envelope of the products of combustion, and 
^ this is sometimes invisible, as in the present 
case, but is sometimes faintly luminous. The 
lower part of the inner cone of the flame is quite cool, and 
consists of unburned gas. Toward the top of the inner cone 
the gas has become heated to a high temperature by the 
burning envelope surrounding it. On reaching the supporter 
of combustion on the outside it is far above its kindling 
temperature, and combustion follows with evolution of much 
heat. The region of combustion just outside the inner cone 
is therefore the hottest part of the flame. 

When the combustible gas is a compound substance, and 
undergoes decomposition before burning, or burns in several 
stages, the number of the cones may be increased, a new cone 
appearing fpr each chemical reaction taking place. In the case 
of ammonia burning in oxygen, the heated ammonia decom- 
poses, thus 2NH3 = N2 + 3II2. The hydrogen then burns to 
form water, and the nitrogen escapes for the most part un- 
changed. There is in this case a double cone surrounding the 
inner cone of unburned gas. When a gaseous hydrocarbon,, 
such as methane,. is burned there are several cones to the flame. 



1 86 Inorganic Chemistry. 

for the methane does not burn as represented in the simple 
equation CH4 + 2O2 == CO2 + 2H2O, but in a series of stages. 
Acetylene is one of the intermediate products and part of the 
luminosity of the gas flame is due to its formation and subse- 
quent combustion. 

Oxidizing and Reducing Flames. Since the tip of 
the outside cone consists of very hot products of combus- 
tion, mixed with oxygen from the air, a substance capable 
■of oxidation placed in this part of the flame is easily oxi- 
dized. The hot gases raise it to a high temperature, and 
under these circumstances it combines with the oxygen 
abundant near the tip of the flame. The oxygen with which 
it combines comes, of course, from the atmosphere, and not 
from the products of combustion. This outer tip of the 
flame is called the oxidizing flame.' 

At the tip of the inner cone the conditions are quite 
different. This region consists of a highly heated combus- 
tible gas, which has not yet reached a supply of oxygen. 
If a substance, rich in oxygen, such as a metallic oxide, is 
placed in this region of the flame, the heated gases combine 
with its oxygen, instead of that of the air, and the substance 
is reduced. This part of the flame is called the reducing 
flame. 

Luminosity of Flames. Much careful study has 
shown that the luminosity of flames is due to a number of 
distinct causes, and that it may therefore be increased or 
diminished in several ways. ( i ) The most obvious of these 
factors is the presence in the flame of incandescent solid 
matter. Thus chalk dust sifted into a non-luminous flame 
renders it luminous. When hydrocarbons form a part of 
the combustible gas, as they do in nearly all illuminating 
gases and oils, some carbon is usually set free in the process 
■of combustion. This is made very hot by the flame, and be- 
comes incandescent, giving out Hght. In a well regulated 
flame it is afterward burned up, and when the supply of oxy- 



Flames — lUuminants. 187 

g-en- is insufficient, it escapes from the flame as lampblack or 
soot. That it is temporarily present in a well burning lumi- 
nous flame may be demonstrated by holding a cold substance, 
as a small evaporating dish, in the flame for a few seconds. 
The cold object cools the carbon below its kindling tempera- 
ture, and it is deposited on the object as soot. 

Substances differ much in the ease with which they be- 
come incandescent in a flame, and in the amount of light which 
they give off when heated to incandescence. The various kinds 
•of mantles for gas lights now so commonly used, are made of 
substances which become incandescent at a low temperature, 
and which then give off an unusual amount of light. These 
mantles are suspended around a non-luminous flame, and it is 
the incandescent substance in the flame which gives the light. 

(2) A second factor in the luminosity of flames is the 
pressure under which the gases are burning. Under in- 
creased pressure there is more matter in a given volume 
of a gas, and the chemical action is more energetic than 
Avhen the gases are rarifled. Consequently there is more 
heat and light. A candle burning on a high mountain gives 
less light than when it burns at the sea level. 

If the gas is diluted with a non-combustible gas, the 
^flfect is the same as to rarify it. There is less combustible 
gas in a given volume. The difference observed between 
burning in air and burning in oxygen is due to the presence 
of so much incombustible nitrogen which serves to dilute 
the oxygen of the air. 

(3) The luminosity also depends upon the temperature 
attained m the combustion. In general the hotter the flame 
the greater the luminosity. Hence cooling the gases before 
combustion diminishes the luminosity of the flame they will 
make, because it diminishes the temperature attained in the 
combustion. 

With these principles in mind, the several kinds of 
flames in common use are readilv understood. 



1 88 Inorganic Chemistry. 

Bunsen Flame. The luminous Bunsen flame is pro- 
duced by hydrocarbons and other gases burning in air. It 
is very hot, and owes its luminosity to its high temperature 
and to the plentiful supply of incandescent carbon in it. When 
air is admitted at the base of the burner, the luminosity dis- 
appears. The cause of this is in part that oxygen is drawn 
up with the gas, and the hot carbon compounds do not have 
to wait for complete combustion until they reach the outer 
cone of the flame. Much less carbon is actually set free 
in any stage of the combustion. The burning gas is also 
diluted and cooled by the air drawn in, which tends to 
diminish luminosity. 

Blowpipe Flame. This is made by blowing into a 
small, luminous Bunsen flame from one side, through a 
blowpipe. This is a tube of the shape shown in 
Fig. 52. The flame is easily directed in any de- 
sired way, and has the oxidizing and reducing 
regions very clearly marked (Fig. 53). It is non- 
luminous from the same causes as render the open 
Bunsen burner non-luminous, the gases from the 
lungs serving to furnish oxygen, and to dilute 
J I the combustible gas. 

p{^''^2. Illuminating and Fuel Gases. A number of 

mixtures of combustible gases consisting largely 

of carbon compounds and hydrogen, find extensive use for 
the production of light and heat. The three chief varieties 

are (i) coal gas, (2) w^ater gas, 
(3) natural gas. The use of ace- 
tylene gas has already been re- 
ferred to. 

Coal Gas is made by heating bi- 
tuminous coal in large retorts out 
p.g 53 of contact with the air. Such ' a 

process is called dry distillation. 




Flames — Illiuninants. - 189 

Soft or bituminous coal contains, in addition to large amounts 
of carbon, considerable quantities of compounds of hydrogen, 
oxygen, nitrogen and sulphur. When distilled, the nitrogen 
is liberated partly in the form of ammonia and partly as free 
nitrogen gas ; the sulphur is converted into hydrogen sul- 
phide ; the oxygen into water, and oxides of carbon. The 
remaining hydrogen is partly set free as hydrogen, and 
partly in combination with carbon in the form of hydro- 
carbons, chiefly as methane, with smaller quantities of many 
others, some of which are liquids and solids at ordinary 
temperatures. The great bulk of the carbon remains be- 
hind as coke and retort carbon. 

The mixture of gases is freed from tarry and solid ma- 
terials by being passed through condensing chambers, and 
from soluble gases such as hydrogen sulphide and ammonia, 
in towers called scrubbers, in which sprays of water are 
played upon the gas. It is then passed over beds of lime or 
iron oxide to absorb carbon dioxide and remaining traces 
of sulphur compounds, and thence into holders for distribu- 
tion. The scrubbers furnish all the ammonia of commerce, 
while from the tarry products, known as coal tar or pitch, 
many valuable products are obtained. Among these are 
benzene, toluene, carbolic acid, naphthalene. 

Water Gas is essentially a mixture of carbon monoxide 
and hydrogen. It is made by passing steam over very hot 
anthracite coal, when the reaction shown in the following 
equation takes place : C + HoO = CO -j- H,. When re- 
quired merely for fuel gas, to produce heat, the gas is at 
once ready for use. When made for illuminating purposes 
it must be enriched, that is illuminants must be added. This 
is accomplished by passing it as it comes from the producer 
into heaters, containing highly heated petroleum oils. The 
gas takes up hydrocarbon gases formed in the decomposi- 



190 ^ Inorganic Chemistry. 

tion of the petroleum oils, which make it burn with a lumi- 
nous flame, whereas the original w^ater gas burns with a non- 
luminous flame of burning carbon monoxide and hydrogen. 
The final composition of enriched water gas does not differ 
very much from that of coal gas, excepting that it has much 
more carbon monoxide in it. 

Natural Gas. This substance, so abundant in many lo- 
calities, varies much in composition, but in a general way 
it resembles coal gas. It contains much more methane, and 
its luminosity is not so great. When used for lighting pur- 
poses, it is usually burned in a burner resembling an open 
Bunsen, the illumination being furnished by an incandes- 
cent mantle. Contrary to statements frequently made, nat- 
ural gas contains no free hydrogen. 

Table Showing Composition of Gases. 

Pennsylvania 

Natural Gas. 

Hydrogen 

Methane 90 . 64 

Illuminants (C2H4) 

Carbon monoxide 

Carbon dioxide 0.30 

Nitrogen 9.06 

Oxygen 

Hydrocarbon vapors 









Enriched 


Coal 


Water 


Water 


Gas. 


Gas. 


Gas. 


41-3 


52.88 


30.00 


43-6 


2 


16 


24.00 


3-9 






12.05 


6.4 


36 


80 


29.00 


2.0 


3 


47 


0.30 


1 .2 


4 


69 


2.50 


0.3 

1-5 






1.50 
1.50 







These are analyses of actual samples, and may be taken 
as about the average for the various kinds of gases. Any one 
of these may vary considerably. The nitrogen and oxygen in 
most cases is due to a slight admixture of air which is difficult 
to exclude entirely in the manufacture and handling of gases. 
' The Electric Furnace. Until recent years about the pnly 
source of heat for practical purposes has been the combustion 
of such substances as coal, the gases just described, and hydro- 
gen. When a very intense heat was desired the oxyhydrogen 



Flames — Illuminants. 



191 




Fig. 54. 



— blow-pipe vv a s 
used. With the 
development o f 
the cheap pro- 
duction of the 
electric current, a 
new source of 
heat, namely, the 
electric furnace, 
is rapidly becom- 
ing an important 
factor in the pro- 
duction of very 
high tempera- 
tures. 
In the simplest type of the electric furnace, the source 
of heat is the arc formed when a strong current passes be- 
tween carbon poles separated a little from each other. The 
substance to be heated is placed in a grapliite crucible, and 
this crucible is put just below the arc in a cavity in some 
difficultly fusible material such as a block of lime. Over this 
block is placed a similar one with a corresponding cavity. The 
arc is therefore formed in a small open space, the walls of 
which consist of difficultly fusible material, which confine the 
heat to a small space and reflect it downward upon the open 
crucible placed below the arc. (See Fig. 54.) 

In some more recent types, the current is passed through 
very infusible material which is packed around the crucible 
containing the substance to be heated, and which becomes in- 
tensely hot when the current is passing through it. This fur- 
nace has the advantage that the temperature reached can be 
more easily controlled. Temperatures estimated to be as high 
as 3500° can be reached in the electric furnace. 



CHAPTER XVIL 

MOLECULAR WEIGHTS, ATOMIC WEIGHTS, FORMULAS. 

Introduction. In the chapter on the Atomic Theory, 
it was pointed out that if it were true that an atom of one 
element always united with a single atom of every element 
with which it forms a compound, and in no other ratio, it 
would be a very easy matter to decide upon figures which 
would represent th-e relative weights of the different atoms. 
It would only be necessary to select some one element as a 
standard, and determine the weight of every element which 
combines with a unit amount (one gram) of the standard 
element. The figures so obtained would evidently repre- 
sent the relative weights of the atoms. 

-For example, if hydrogen is selected as the standard 
element, it will be found that i gram of hydrogen combines 
with 7.94 grams of oxygen, forming 8.94 grams of water. 
Any weight of water no matter how small will contain oxy- 
gen and hydrogen in the ratio i : 7.94. Consequently the 
molecule of water will contain the elements in this ratio, and 
if a single atom of each is present, these figures must repre- 
sent their relative weights. 

But the law of multiple proportion at once reminds us 
that two elements may unite in several proportions ; and 
there is no simple way to decide in which one of the com- 
pounds so formed the molecules consist of a single atom 
of each of the elements present. Consequently the problem 
of deciding upon the relative atomic weights is not an easy 
one. To the solution of this problem we must now turn. 

When Dalton first advanced the atomic theory he at- 
tempted to solve this problem by very simple methods. He 
thought that when ^only one compound of two elements is 
known, it is reasonable to suppose that it contains one atom 

(192) 



Molecular Weights, Atomic Weights, Formulas. 193 

of each element. He therefore gave the formula HO to 
water, and HN to ammonia, CH to ethylene gas (C2H4). 
When more than two compounds were known he assumed 
that the most familiar or the most stable compound had the 
simple formula. He then determined the atomic weight as 
explained above. The results he obtained were contradictory, 
and ver}' far from satisfactory, and it was soon seen that some 
other wa}', resting on much more scientific grounds, must be 
found to decide what compounds have a single atom of each 
element present. 

Three distinct steps are involved in the selection of 
the proper number to express the atomic weight of an ele- 
ment : 

I. Determination of the Equivalent. 

Definition. By the equivalent of an element is meant 
the mass of the element which will combine with a fixed mass 
of some other elem-cnt chosen as a standard. Since hvdro- 
gen has the smallest equivalent, it is chosen as the standard ; 
the equivalent of an element is therefore the amount of it 
wdiich combines with a gram of hydrogen, or displaces a 
gram of hydrogen from an acid. Thus, i gram of hydrogen 
combines with 7.94 grams of oxygen, 15.92 grams of sul- 
phur, 35.18 grams of chlorine. 32.45 grams of zinc and 
31.92 grams of copper displace a gram of hydrogen from 
acids. These figures therefore represent the equivalent 
weights of these elements. 

Indirect Determination of the Equivalents. Many of the 
elements do not form compounds with hydrogen which can 
he easily obtained and analyzed, nor do they all displace 
hydrogen from acids in such a way that the hydrogen can 
be collected and measured. Nearly all the elements do, how- 
ever, combine with oxygen, and since the equivalent of oxy- 
gen is 7.94 the equivalent of an element can often be deter- 
mined hy finding out how much of it combines with 7.94 
•grams of oxygen. 



194 Inorganic Chemistry. 

Relation of Equivalents to Atomic Weights. AccordmsT 
to the atomic theory combination always takes place between 
whole numbers of atoms. One atom unites vvith one other, 
or with two or three ; or two atoms may unite with three, or 
three with five. When oxygen combines with hydrogen the 
combination must be between definite numbers of the two 
kinds of atoms. If one atom of hydrogen combines with one 
of oxygen, and we find by experiment that the ratio by 
weight is i gram of hydrogen to 7.94 grams of oxygen, 
then the atom of oxygen must be 7.94 times the weight of 
the hydrogen atom. If two atoms of hydrogen combine 
with one of oxygen, the oxygen atom is 15.88 times as 
heavy as the hydrogen atom. 

It is evident that the equivalent must bear a simple re- 
lation to the real atomic weight ; it is either some multiple 
or aliquot part of it. And since the equivalents can be very 
accurately determined, we can in this way obtain a value 
for each element which stands in this very simple relation 
to its atomic weight. 

Combining Weights. The equivalent of an element de- 
notes not only the amount of it which will combine with 
a gram of hydrogen, but also the amount which will 
combine with the equivalent weight of another element. We 
have seen that the figures 7.94, 15.92 and 35 18 are the 
equivalents respectively of oxygen, sulphur and chlorine, 
and that 32.45 is the equivalent of zinc. Now it has been 
found by experiment that 32.45 grams of zinc will combine 
with 7.94 grams of oxygen, 15.92 grams of sulphur and 
35.18 grams of chlorine. These figures may therefore be 
called the combining weights of the elements named. It is 
important to notice that the fact that a number can be found 
for each element which indicates the proportion in which 
the element combines with other elements is not a part of a. 



Molecular JVeiglifs, Atomic Weights, Formulas. 195 

theory, but is a fact of nature, just as the fact expressed by 
the law of muhiple proportion is a fact and not a theory. 

Elements with more than one Equivalent. It will be re- 
membered that oxygen combines with hydrogen in two ra- 
tios. In the one case i gram of hydrogen combines with 
7.94 grams of oxygen to form water ; in the other i gram 
of hydrogen combines with 15.88 grams of oxygen to form 
hydrogen dioxide. The equivalents of oxygen are therefore 
7.94 and 15.88. Barium combines with oxygen in two pro- 
portions : in barium oxide the proportion is 7.94 grams of 
oxygen to 68.72 grams of barium ; in barium dioxide the 
proportion is 7.94 grams of oxygen to 34.36 grams of ba- 
rium. 

In each case the one equivalent is a simple multiple of 
the other, so the fact that there may be two equivalents does 
not add uncertainty to the case ; for all we knew before 
was that the true atomic weight is some multiple of the 
equivalent. 

2. The Determination of Molecular Weights. 
To decide the question as to which multiple of the equivalent 
correctly represents the atomic weight of an element, it has 
been found necessary to devise a method of determining 
the molecular weights of compounds containing the eleirient 
in question. Since the molecular weight of a compound is 
merely the sum of all the weights of all the atoms present 
in it, it would seem to be impossible to determine the mo- 
lecular weight of a compound without first knowing the 
atomic weights of the constituent atoms, and how many 
atoms of each element are present in the molecule. But 
certain facts have been discovered which suggest a way in 
which this can be done. 

Simplicity of the Gas Lazvs. We have already found 
that the laws which govern the. conduct of gases are very 
simple and seem to apply to all gases whatever their chemical 



19^ Inorganic Chemistry. 

properties may be. Boyle's Law tells us that a given vol- 
ume of any gas under standard conditions is affected in the 
same degree by a given change in pressure. Dalton's Law 
states the fact that a given volume of any gas under stand- 
ard conditions is affected in the same degree by a given 
change in temperature. Gay-Lussac's Law shows that very 
simple relations hold between the volumes of combining 
gases, as well as between them and the volume of the 
gaseous product formed. It is evident therefore that the 
structure of gases must be very simple, and that it is much 
the same in all gases. 

Avogadro's Hypothesis. In 1811, Avogadro, an Italian 
physicist, suggested that if we assume that all gases, under 
the same conditions of temperature and pressure, have the 
same number of molecules in a given volume, we will have 
a probable explanation of these simple relations. Plainly we 
cannot prove such an hypothesis, but there are a great 'many 
facts known which make it seem p)robable that it expresses 
the truth. Avogadro's hypothesis then is : — Equal volumes 
of all gases, under the same conditions of temperature and 
pressure, contain the same number of molecules. 

Avogadro's Hypothesis and Molecular Weights. As- 
suming that Avogadro's hypothesis is correct, we have a 
very simple means for deciding upon the relative weights of 
molecules. For if equal volumes of two gases contain the 
same number of molecules, the weights of the two volumes 
must be in the same ratio as the weights of the indivdual 
molecules which they contain. If we adopt some one gas 
as a standard we can express the weights of all other gases 
as compared with this one, and the same figures will express 
the relative weights of the molecules of which the gases are 
■composed. 

Hydrogen as Standard. Hydrogen, being the lightest 
■of all gases, is a suitable standard. On comparing the 
weights of equal volumes of hydrogen and other gases, it is 




Molecular Weights, Atomic Weights, Formulas. 197 

found that oxygen is 15.88 times as heavy as hydrogen: 
nitrogen 14 times as heavy; chlorine 35.18 and hydrochloric 
acid 18.09 times as heavy. Since these equal volumes have 
the same number of molecules in them, it follows that these 
figures express the weights of the molecules of the sub- 
stances named, compared with the hydrogen molecule as 
unity. 

The relative weights of equal volumes of two gases can be 
easily determined. A small flask such as is shown in the Fig. 55 
is filled with one of the gases, and after the 
temperature and pressure have been noted, the 
flask is sealed up and weighed. The top is 
then broken off of the sealed end, the flask filled 
with the second gas, and its weight determined. 
If the weight of the empty flask is subtracted 
P^g'- 5^- from these two weighings, the relative weights 

of the gases is readily found. 

The Hydrogen Molecule Contains Tzi'o Atoms. It has 
already been shown that when hydrogen and chlorine com- 
bine, one volume of hydrogen combines with one volume of 
chlorine to form two volumes of hydrochloric acid gas. Let 
us suppose that the equal volumes of hydrogen and chlorine 
contained 1,000 molecules; then by Avogadro's law, the 
double volume of hydrochloric acid formed will contain 
2,000 molecules. But each of these 2,000 molecules must 
contain at least one atom of hydrogen, or 2,000 atoms :n all^ 
and the 2,000 hydrogen atoms came from 1,000 molecules of 
hydrogen. It follows therefore that the hydrogen molecule 
consists of at least tv/o atoms. Evidently this reasoning 
only proves that there are at least two atoms in the hydro- 
gen molecule ; there may be more than that, but as there 
is no evidence of any kind that this is the case we assume 
that the molecule contains only two atom.s. 

We have just seen that the oxygen molecule is 16 times 
as heavy as the hydrogen molecule ; and if the hydrogen 
molecule contains two atoms, the oxygen molecule is evi- 



198 Inorganic Chemistry. 

dently ^2 times as heavy as the hydrogen atom. For many 
reasons the hydrogen atom is a more convenient standard 
than the hydrogen molecule with which to measure the rela- 
tive weights of other molecules. Consequently to determine 
the molecular weight of a gas we find its weight compared 
with the weight of an equal volume of hydrogen gas, and 
multiply this value by two; hence molecular weight := 2X 
density (compared with hydrogen). In determining the 
density of gases it is customary to use air as a standard 
rather than hydrogen. Since air is 14.44 times as heavy as 
hydrogen, we can substitute the density compared with air 
in the above equation instead of the density compared' with 
hydrogen, if at the same time we multiply by 14.44. We 
have then the formula ; molecular weight = 2 X 14-44 X 
density (referred to air) or in brief: M = d X 28.88. 

3. Deduction of Atomic Weights from Molecular 
Weights and Equivalents. We have now seen hOw the 
equivalent of an element, which is some aliquot part of the 
atomic weight, and the molecular weight of compounds con- 
taining the element, can be obtained. Let us see how it is 
possible to decide which multiple of the equivalent really is 
the true atomic weight. An example will make the reason- 
ing clear, so let us suppose that the equivalent of oxygen 
has been found to be 7.94 and that it is desired to obtain its 
atomic weight. The next step is to obtain the molecular 
weights of a large number of compounds containing oxygen. 
The following will serve : 

Water 

Hydrogen dioxide 

Carbon monoxide 

Carbon dioxide 

Nitro2:en peroxide 

Phosphorus pentoxide . . 



Molecular 

weight 
(in round 
numbers). 

18 


Part due 
to oxygen 
(in round 
numbers). 

16 


34 


32 


28 


^6 


44 


32 


46 


3^- 


142 


80 



Molecular JVciglits, Atomic JVeigJits, Foniiulas. 199 

The figures under molecular weight gives the weight of 
the molecules compared with the hydrogen atom, and have 
been determined as just described. The figures in the sec- 
ond colimm show what part of this total weight is due to 
the oxygen in the compound, and have been found by anal- 
ysis of the substances. Now if the molecular weights repre- 
sent the sum of the atomic weights, then the figures in the 
second column must also represent the sum of the weights 
of the oxygen atoms in the molecules ; and since the mole- 
cule is composed of whole atoms, the figures representing 
the wei.:^"hts of the atoms must be the sum of the weights of 
Avhole atoms of oxygen. Now among all the compounds 
containing oxygen, there will doubtless be some containing 
only one atom of oxygen, and so we should find some small 
fxgure in the column expressing the weights of the sum of 
the oxygen atoms which is the weight of a single oxygen 
atom, and all the other weights should be multiples of this 
least weight. No compound has ever been found which con- 
tains less than 16 parts of oxygen in the molecular weight 
of a compound,' and all other weights are multiples of 16. 
\A'e tlierefore assume that 16 is the atomic weight of oxygen. 

Accurate Deteruiivcation of Atomic Weights. ]\Iolecu- 
lar weic;hts cannot be determined very accurately, and con- 
sequentlv the part in the molecular weights due to oxygen is 
a little uncertain. All we can tell by this method is that the 
true weight is very near 16. The equivalent can however be 
determined very accurately, and we have seen that it is some 
aliquot part of the true atomic weight. Since molecular 
weight determination has shown that in the case of oxygen 
the atomic weight is near 16, and we have found the equiva- 
lent to be 7.94, it is evident that the true atomic weight is 
twice the equivalent, or 7.94 X 2 = 15.88. 

These then are the steps necessary to establish the 
atomic Vv'eight of an element: 



200 Inorganic, Chemistry. 

1. Determine the equivalent accurately by analysis. 

2. Determine the molecular weight of a large number 
of compounds of the element, and by analysis the part of 
the molecular weight due to the element. The smallest num- 
ber so obtained will be approximately the atomic weight. 

3. Multiply the equivalent by the small whole num- 
ber (usually I, 2 or 3), which will make a number very 
close to the approximate atomic weight. The figure so ob- 
tained will be the true atomic weight. 

Other Methods of Determining Molecular PV eights. It 
will be noticed that the reason Avogadro's Law furnishes 
us with a method by which we can determine the relative 
weights of the molecules of two gases is that it enables us 
to tell when we are dealing with an equal number of the twa 
kinds of molecules. If by any other means we can get this 
information, we can make use of the knowledge so gained to 
determine the molecular weights of the two substances. 

Raoulfs Laws. Two laws have been discovered which 
give us just such information. They are known as Raoult's 
Laws, and can be stated as follows : 

1. When weights of substances which are proportional 
to their molecular weights are dissolved in the same weight 
of solvent, the rise of the boiling point is the same in each 
case. 

2. When weights of substances which are proportional 
to their molecular weights are dissolved in the same weight 
of solvent, the lowering of the freezing point is tlie 'same 
in each case. 

By taking advantage of these laws it is possible to de- 
termine when two solutions contain the same number of 
molecules of two dissolved substances, and consequently the 
relative molecular weights of the two substances. 

Direct Method of Determining Atomic Weights. In 
1819 Dulong and Petit discovered a very interesting rela- 



Molecular Weights, Atomic Weights, Formulas. 201 

tion between the atomic weight of an element and its specific 
heat, which holds true for elements in solid state. If equal 
weights of two solids, say lead and silver, are heated through 
the same range of temperature, as from 10° — 20°, it is found 
that very different amounts of heat are required. The 
amount of heat required to change the temperature of a 
sohd or a liquid by a definite amount compared with the 
amount required to change the temperature of an equal 
weight of water by the same amount is called its specific 
heat. Dulong and Petit found that the specific heat of an 
element in the solid form multiplied by its atomic weight is 
the same in all cases, and is equal to a nearly constant value 
of 6.25. That is atomic weight X specific heat — 6.25 

and consequently atomic weight = r^ — r — , This law 

^ -^ ^ specific heat 

is not very accurate, but it is often found possible by means 

of it to decide upon what multiple of the equivalent is 

the real atomic weight. Thus the specific heat 61 iron is 

found to be 0.112 and its equivalent is 28. 6.25 -i- 0.112 := 

55.8. We see therefore that the atomic weight is twice the 

equivalent or 56. 

How Formulas are Determined. It will be well in 
connection with molecular weights to consider how the for- 
mula of a compound is decided upon, for the two subjects 
are very closely associated. Some examples will make the 
method followed clear. 

The molecular weight of a substance containing hydro- 
gen and chlorine was 36.4. By analysis 36.4 parts of the sub- 
stance was found to contain i part of hydrogen and 35.4 
parts of chlorine. As these are the simple atomic weights 
of the two elements, the formula and the compound must 
be HCl. 

A substance consisting of oxygen and hydrogen was 
found to have a molecular weight of 34. Analysis showed 



^02 Inorganic Chemistry. 

that in 34 parts of the substance there are 2 parts of hydro- 
gen and 32 parts of Oxygen. Dividing these figures by the 
atomic weights of the two elements we get 2-1-1 = 2 for 
H : 32 ^ 16 = 2 for O. The formula is therefore HoO^. 

A substance containing 2.04% H, 32.6% S and 65.3% 
O was found to have a molecular weight of 98. In this 98 
parts of the substance there are 98 X 2.04% = 2 parts of H, 
98X32.6% = 32, S and 98X65.3% =64, O. If the 
molecule weighs 98, the hydrogen atoms present must to- 
gether weigh 2, the sulphur atoms 32, and the oxygen atoms 
64. Dividing these figures by the respective atomic weights 
of the three elements we have for H, 2-^1 = 2 atoms, for 
S, 32 -f- 32 = I atom, for O, 64 -^ 16 = 4 atoms. Hence 
the formula is H0SO4. 

We have then this general procedure : Find the per- 
centage composition of the substance and also its molecular 
weight. Multiply the molecular weight in succession by the 
percentage of each element present to find the amount of 
the element in the molecular weight of the compound. The 
figures so obtained will be the respective parts of the mo- 
lecular weight due to the several atoms. Divide by the 
atomic weight of the elements, and the quotient will be the 
number of atoms present. 



CHAPTER XVIII. 

THE NITROGEN FAMILY. 

Tlie members of the Nitrogen Family given in the order 
of their atomic weights are nitrogen, phosphorus, arsenic, 
antimony and bismuth. Nitrogen has been studied in a former 
chapter. The remaining mem.bers of the family will new be 
discussed. 

PHOSPHORUS. 

History. The element phosphorus was discovered 
by the Alchemist Brandt of Hamburg in 1669, while searcli- 
ing for the philosopher's stone. Ovv^ing to its peculiar prop- 
erties, and the secrecy which was maintained about its prepa- 
ration, it remained a very rare and costly substance, until 
the demand for it in the manufacture of matches brought 
about its production on a large scale. 

Occurrence. Owing to its great chemical activity, 
phosphorus never occurs free in nature. In the form of 
phosphates it is very abundant and widely distributed. Phos- 
phorite and sombrerite are mineral forms of calcium phos- 
phate, while apatite consists of calcium phosphate together 
with calcium chloride or fluoride. These minerals form 
very large deposits, and are extensively mined for use as fer- 
tilizers. Calcium phosphate is a constituent of all fertile soil, 
having h-een supplied to the soil by the disintegration of 
rocks containing it. It is the chief mineral constituent of 
bones of animals, and bone ash is therefore nearly pure cal- 
cium phosphate. 

Preparation. Phosphorus is now manufactured from 
bone ash or a pure mineral phosphate, by heating the phos- 
phate with sand and carbon in an electric furnace. 

To understand the . reaction which occurs, it . must be 
remembered that a volatile acid anhydride is expelled from 

(203) 



204 Inorganic Chemistry. 

its salts when heated with an anhydride which is not volatile. 
Thus when sodium carbonate and silicon dioxide are heated 
together, the following reaction takes place: 
Na2C03+Si02=Na2Si03+C02. 
Silicon dioxide is a less volatile anhydride than phosphoric 
anhydride, P2O5, and when strongly heated with a phosphate, 
the phosphoric anhydride is driven out, thus: 

Ca3(P04)2+3Si02=3CaSi03+P205. 
If carbon is added before the heat is applied, the P2O5 is re- 
duced to phosphorus at the same time, according to the equation 

P205+5C=2P+sCO. 
The phosphorus vapor, as it escapes from the furnace, is con- 
ducted under water, where it is condensed into solid form. 

In the older process, which is still in use, bone ash is 
treated with sulphuric acid, forming calcium sulphate and 
phosphoric acid. The phosphoric acid is evaporated to a 
syrup and mixed with sawdust or carbon. The mixture is 
then dried, and finally heated intensely in an earthen retort, 
when phosphorus is set free according to the equation : 

2 H3PO, + 5 C = 3 H,0 + 2 P + 5 CO. 
The phosphorus so prepared contains many impurities. 
These ^ are removed in various ways, usually by melting the 
crude phosphorus in warm water, and pressing it through 
filters of bone black or chamois skin. The product is finally 
distilled. 

Physical Properties. The purified phosphorus is a 
pale, yellowish, translucent, waxy solid which melts at 45% 
and boils at 290°. It melts therefore in warm water, and 
can be cast into any convenient form under water. It is 
usually sold in the market in the form of sticks. It is quite 
soft, and can be easily cut with a knife. This must always 
be done while the phosphorus is covered with water, since 
the element is extremely inflammable, and the friction ef 
the knife blade is almost sure to set it on fire if cut iri the 
air. It is not soluble in water, but is freely soluble in some 
other liquids, notably in carbon disulphide. Its density is 1.8. 



The Nitrogen Family. 205 

Chemical Properties. Exposed to the air, phos- 
phorus slowly combines with oxygen, and in doing so, emits 
a pale light or phosphorescence, which can be seen only in a 
da^k place. The heat of the room may easily raise the tem- 
perature to the kindling point, when phosphorus burns with 
a sputtering flame, giving off dense fumes of oxide of phos- 
phorus. 

The heat of the body is sufficient to raise phosphorus 
above its kindling temperature, and for this reason phos- 
phorus should always be handled with forceps and never 
w^ith the bare fingers. Burns occasioned by it are very pain- 
ful and slow in healing. Phosphorus burns brilliantly in 
oxygen, and combines directly wuth many other elements, 
especially with sulphur and the halogens. 
P + 3 Cl = PCI3. 

Phosphorus is very poisonous, 0.2 — 0.3 gram being a 
fatal dose. Ground up with flour and water or similar sub- 
stances, it is often used as a poison for rats and other vermin. 

Red Phosphorus. When phosphorus is heated for 
some time at a temperature of about 250° — 300°, it under- 
goes a remarkable change. It loses its transparency and is 
converted into a dark, chocolate colored, opaque powder hav- 
ing a density of about 2.1. It no longer takes fire easily, 
and can be handled in dry form like any other substance. 
It is not poisonous, and in fact seems to be an entirely dif- 
ferent substance. When distilled, and quickly condensed, 
it changes back again into ordinary yellow phosphorus. Still 
another, or black modification of phosphorus is known, which 
is somewhat metallic in appearance. These allotropic forms 
recall the similar modifications of sulphur and carbon. 

Uses of Phosphorus. The chief use of phosphorus 
is in the manufacture rf matches. Common matches are 
made by first dipping the match sticks into some inflammable 
substance such as melted paraffine, and afterward into a 



2o6 



Inorganic Chemistry 



paste consisting of (i) phosphorus, (2) some oxidizing sub- 
stance such as manganese dioxide, or potassium chlorate, and 
(3) a binding material, usually some kind of glue. On fric- 
tion the phosphorus is ignited, the combustion being sus- 
tained by the oxidizing agent, and communicated to the wood 
by the burning paraffine. 

In sulphur matches the paraffine is replaced by sulphur. 
Jn safety matches, red phosphorus, an oxidizing agent, and some 
gritty material such as emery is placed on a strip on the side 
of the box, while the match tip is provided as before with an 
oxidizing agent and an easily oxidized substance, usually anti- 
mony sulphide. The match cannot be easily ignited by fric- 
tion, save on the prepared surface. 

Compounds of Phosphorus with Hydrogen. Phos- 
phorus forms several compounds with hydrogen, the best 
known of which is phosphine, PHg, analgous to ammonia, 

NH3. 

Preparation of Phosphine. Phosphine is usually made 
by heating phosphorus with a strong solution of potassium 
hydroxide, the reaction being a complicated 
one. The experiment can be conveniently 
made in the apparatus shown in Fig. 56. 
A strong solution of potassium hydroxide 
)gether with several small bits of phos- 
phorus are placed in the flask A, and a 
current of coal gas is passed 
into the flask through the 
tube B, until all the air has 
been displaced. The gas is 
then turned off, and the flask 
is heated. Phosphine is 
formed in small quantities, 
and escapes through the de- 
livery tube, the exit of which 




c:;r;:ri:;;i> 






Fig. 56. 



The Nitrogen Family. 207 

is just covered by tlie water in the vessel C. Each bubble of 
the gas, as it escapes into the air, takes fire, and the product 
of combustion forms beautiful small rings, which float un- 
broken for a considerable time in quiet air. 

Properties. Phosphine is a gas of unpleasant odor, 
and is exceedingly poisonous. Prepared in the way just 
described, it is spontaneously inflammable, but this is due to 
impurities. It has much more feeble basic properties than 
has ammonia, forming unstable salts with the halogens acids,, 
but with no others. 

PH3 + HCl = PH.Cl 
The product of this reaction is called phosphonium chloride, 
and is analogous to ammonium chloride. 

Oxides of Phosphorus. Phosphorus forms two well 
known oxides : the trioxide, P0O3, and the pentoxide, PoOg, 
sometimes called phosphoric anhydride. When phosphorus 
burns in an insufficient supply of air the product is par- 
tially the trioxide : in oxygen or free access of air the pent- 
oxide is formed. ' The pentoxide is much the better known 
of the two. It is a snow white, voluminous powder, whose 
most marked property is its great attraction for water. It 
has no chemical action upon most gases, so that they can be 
very thoroughly dried by allowing them to pass through 
properly arranged vessels containing phosphorus pentoxide. 

Acids of Phosphorus. Both of the oxides of phos- 
phorus are acid anhydrides : that is they combine with water 
to form acids. Unlike the corresponding nitrogen anhy- 
drides, they show a marked tendency to form acids contain- 
ing three hydrogen atoms, so that the relation between 
acid and anhydride is not so simple as with the nitrogen 
acids. Thus the trioxide combines with water to form the 
tribasic phosphorous acid : 

P.O3 + 3 H3O = 2 H3PO3. 



2o8 Inorganic Chemistry. 

This acid, as well as its salts, is not very often met with. 
Its most interesting property is its tendency to take up oxy- 
gen and pass over to phosphoric acid : 
H3PO3 + O = H3PO,. 
It is therefore a good reducing agent. 

Phosphoric Acids. Phosphorus pentoxide is able 
to combine with water in three proportions to form three 
acids, the names and formulas of which are : 

HPO3 metaphosphoric acid 

H3PO4 orthophosphoric acid 

H4P2O7 pyrophosphoric acid. 

(i) Metaphosphoric Acid. This acid is formed when 
phosphorus pentoxide dissolves in water 
P2O, + HoO = 2 HPO3 
It is a crystalline solid, and is very stable towards heat, so 
that it can be fused and even volatilized without decomposi- 
tion. On cooling from the fused state it forms a glassy 
solid, and the acid is on this account often called glacial 
phosphoric acid. It possesses the property of dissolv 
small quantities of metallic oxides, and acquiring chaiu.^,- 
teristic colors from them. On this account it is often used to 
detect the presence of such metals in a compound as give 
these colors with metaphosphoric acid. 

(2) Orthophosphoric Acid. When a solution of 
metaphosphoric acid is boiled, the acid combines with an 
additional amount of water, forming phosphoric acid, having 
the formula HgPO^. 

HPO3 + H2O = H3PO,. 
Phosphoric acid, sometimes called orthophosphoric acid, 
can be obtained in the form of large colorless crystals, 
by evaporating its solutions. When strongly heated,, it 
loses water and metaphosphoric acid is formed again. 
It is not volatile, and will therefore drive most of the 



The Nitrogen Family. 209 

common acids from their salts when the salt and phos- 
phoric acid are heated together. 

Phosphoric acid is a tribasic acid ; that is all three of 
its hydrogen atoms have acid properties. It can therefore 
form two acid salts and a normal salt, \Hien acted upon by 
bases. 

NaOH + H3PO4 r= NaH^PO^ + HoO. (primary acid phos- 
phate.) 

2 NaOH + H3PO4 = NaH^PO^ + 2 H^O. (secondary acid 
phosphate.) 

3 NaOH + H3PO4 =: NagPO, + 3 H^O. (tertiary acid 
phosphate.) 

it can easily form mixed salts, that is salts containing 
two different metals. Microcosmic salt is such a mixed salt 
^^ sodium and ammonium: it has the formula, Na(NH4) 
HPO4, H2O. 

Ortho phosphates. The orthophosphates form an im- 
portant class of salts. The normal, or tertiary, phosphates 
are nearly all insoluble, and as we have seen, occur in large 
^mtities in nature. The primary phosphates are for the 
ill^:.■st part soluble in water. 

(3) Pyrophosphoric Acid. On heating a primary 
phosphate, it loses oxygen and hydrogen in the form of 
iter, and a metaphosphate is formed. 

NaH.PO, r= NaPOg + H,0. 

On heating a secondary phosphate, water is eliminated 
form two molecules of the salt. 

2 Na^HPO, = Na^P.O, -f H.O. 

The product of this reaction is the sodium salt of the 
acid H^PoO-, called, from its method of preparation, pyro- 
phosphoric acid. 

The relation of these four acids to the two anhydrides 
can be readily seen from the following equations : 



210 Inorganic Chemistry. ' 

" P2O3 + 3 H.O 3= -2 H3PO3. - phosphorous acid. 

PoO- 4- I HoO = 2 HPO3. m^taphosphoric acid. 

P2O5 + 2 HoO = H4P2O7. pyrophosphoric acid. 

P2O5 + 3 HoO = 2 H^PO^. orthophosphoric acid. 

Fertilizers. When crops are produced year after 
year on the same field certain constituents of the soil essen- 
tial to plant growth are removed, and the soil becomes im- 
poverished and unproductive. To make the land once more 
fertile, these constituents must be replaced. The calcium 
phosphate of the mineral deposits or of bone ash serves well 
as a material for restoring phosphorus to soils exhausted 
of that essential element; but a more soluble substance, 
which the plants can more readily assimilate, is desirable. 
It is better therefore to convert the insoluble calcium phos- 
phate into the soluble primary phosphate,.before it is applied 
as fertilizer. This is commonly done by treatment of the 
normal calcium phosphate with the amount of sulphuric 
acid required to convert it into the primary phosphate, in 
accordance with the equation 

Ca3(PO,)_o + 2 H^SO, = Ca(HoPOj2 + 2 CaSO,. 
The resulting mixture is a powder, which is sold as "super- 
phosphate of lime." 

ARSENIC. 

Occurrence. Arsenic occurs in considerable quanti- 
ties in nature as the native element, as the suphides realgar, 
AS2S0 and orpiment, As.,So, as oxide, AS2O3, and as a con- 
stituent of many metallic sulphides such as arsenopyrite, 
FeAsS. ' . 

Preparation. The element is prepared by purifying 
the native arsenic, or by heating the arsenopyrite in iron 
tubes, w^hen the following reaction occurs: 
FeAsS = FeS + As. 



The Nitrogen Family. 211 

The arsenic being volatile, condenses in chambers con- 
nected with the heated tubes. It is also made from the 
oxide by reduction with carbon : 

2 As.Oa + 3 C = 4 As + 3 CO2. 

Properties. Arsenic is a steel gray, metallic looking 
substance, of density 5.73. Though resembling metals in 
appearance, it is quite brittle, being easily powdered in a mor- 
tar. When strongly heated it sublimes, that is it passes into 
a vapor without melting, and condenses again to a crystalline 
solid when the vapor is cooled. Like phosphorus it can be 
obtained in several allotropic forms. It alloys readily with 
some of the metals, and finds its chief use as an alloy with 
lead which is used for making shot, the alloy being harder 
than pure lead. Unlike most of its compounds it is not 
poisonous. 

Arsine. AVhen any compound containing arsenic is 
brought into the presence of nascent hydrogen, arsine, AsHg, 
corresponding to phosphine and ammonia, is formed. The 
reaction when oxide of arsenic is so treated is 
As.Og + 12 H = 2 AsHg + 3 H2O. 

Arsine is a gas with a peculiar garlic-like odor, and is 
intensely poisonous. A single bubble of the pure gas has 
been known to prove fatal. It is an unstable compound, 
decomposing into its elements when heated to a moderate 
temperature. It is combustible, burning with a pale bluish- 
white flame to form arsenic trioxide and water when air 
is in excess : 

2 AsHg + 6 O = AS2O3 + 3 H2O. 

When the supply of air is deficient, water and metallic 
arsenic are formed : 

2 AsHg + 30 = 3 H2O -f 2 As. 

These reactions make the detection of even minute quan- 
tities of arsenic a very easy problem. 



212 



Inorganic Chemistry. 



Marsh's Test for Arsenic. The method devised by Marsh 
is the one most frequently employed. The form of apparatus 
iised is shown in Fig. 57. Hydrogen is prepared in the genera- 
tor A by the action of dilute sulphuric acid on zinc and is dried 
by passing through the tube B filled with calcium chloride. 
Jt is then conducted through the hard glass tube C and after 







K^ 



i l> A-.V. ■^.'Ly 



^ 



^ 



Fig. 57. 

all the air is expelled from the apparatus, is ignited at the 
jet D. The substance suspected to contain the arsenic is dis- 
solved in dilute sulphuric acid and slowly added to the hydro- 
gen generator through the funnel tube. Any arsenic present 
combines with the hydrogen to form arsine, which being a 
gas, escapes from the generator with the hydrogen. If the tube 
C is kept quite hot at some point near the middle, the arsine is 
decomposed at this point, the' arsenic being deposited in the 
cooler portions of the tube just beyond the flame, in the form 
of a shining black mirror. If the tube is not heated the a-rsine 
passes along with the hydrogen and burns at the jet. If a 
small porcelain evaporating dish be crowded down into the 
flame, the temperature of the flame is reduced below the 
kindling temperature of arsenic, which is deposited on the 
white porcelain as a black spot. After the dish becomes heated 
the arsenic is no longer deposited: hence it is well to have the 
dish filled with cold water. 



The Nitrogen Family. 213 

Oxides of Arsenic. Arsenic forms two oxides, AsoOg 
and AsgOg, corresponding to those of phosphorus. Of these 
arsenious oxide, AsgOg, is much better known, and is the 
substance usually called white arsenic or merely arsenic. It 
is found as a mineral, but is usually obtained as a by- 
product in burning pyrite in the sulphuric acid industry. 
The pyrite has a small amount of arsenopyrite in it, and 
when this is burned, arsenic trioxide is formed together 
with sulphur dioxide : 

2 FeAsS + 10 O = Fe.Og + As^Og + 2 SO^. 
The arsenious oxide is condensed in appropriate chambers. 
It is a rather heavy substance, obtained either as a crystalline 
powder, or as large, vitreous lumps, recalling lumps of porce- 
lain in appearance. It is very poisonous, 0.2 — 0.3 gram being 
a fatal dose. It is used largely as a poison, since it is nearly 
tasteless, and does not act very rapidly. This slow action 
is due to the fact that arsenic is not very soluble, and is very 
slowly absorbed by the system. It is also used as a chemical 
reagent in glass making and in the dye industry. 

Acids of Arsenic. Like the corresponding oxides of 
phosphorus, the oxides of arsenic are acid anhydrides. In 
solution they combine with bases to form salts, correspond- 
ing to the salts of the acids of phosphorus. Thus wc have 
salts of the following acids 

H3ASO3 arsenious acid. 

HAsOg metarsenic acid. 

H3ASO4 orthoarsenic acid. 

H4AS0O- pyroarsenic acid. 

Several other acids of arsenic are also known. Not all 
of. these can be obtained as free acids, since they tend to 
lose water and form the oxides. Thus instead of obtaining 
the acid H^AsOo the oxide As^Oo is obtained. 
2 H3ASO3 = As,03 + 3 H,0. 



214 Inorganic Chemistry. 

Salts of all the acids are known, however, and some of 
them have commercial value. Most of them are insoluble, 
and some of the copper salts which are green, are used as 
pigments. Paris green, which has a complicated formula, 
is a well known insecticide. 

The most efficient antidote for arsenic poisoning is fer- 
ric hydroxide. This converts the soluble arsenious oxide 
into iron arsenite, which being insoluble, is not poisonous. 
AS2O3 + 2 Fe(OH)3 = 2 FeAsOg + 3 H^O. 

Magnesium hydroxide is administered at the same time 
to neutralize the acid naturally present in the stomach. 

Sulphides of Arsenic. When hydrogen sulphide is 
passed into an acidified solution containing an arsenic com- 
pound, the arsenic is precipitated as a bright yellow sul- 
phide. Thus, 

2 H3ASO, + 3 H,S = As.Sg + 3 H^O. 
2 H.AsO, + 5 H^S = As.S, -f 5 H2O. 

In this respect, arsenic resembles the metallic elements, 
many of which produce sulphides under similar conditions. 
The sulphides of arsenic, both those produced artificially, 
and those found in nature, are used as yellow pigments. 

ANTIMONY. 

Occurrence. It occurs in nature chiefly as the sul- 
phide, Sb^Sa, called stibnite, though it is also found as oxide, 
and as a constituent of many complex minerals. It is a 
rather widely distributed and abundant element. 

Preparation. Antimony is prepared from the sul- 
phide in a very simple manner. The sulphide is melted 
with scrap iron in a furnace, the iron combining with the 
sulphur to form a slag or liquid layer of melted iron sul- 
phide, while the heavier antimony settles to the bottom, and 



The Nitrogen Family. 215 

IS drawn off from time to time. The reaction involved is 
represented b}- the equation : 

SKSe + 3 Fe = 2 Sb + 3 FeS. 

Physical Properties. Antimony is a bluish white, 
metallic-looking substance whose density is 6.7. It is 
highly crystalline, hard and very brittle. It has a rather 
low melting point (432°) and expands very noticeably on 
solidifying. 

Alloys. Some metals when melted together thoroughly 
intermix, and on cooling form a homogeneous, me- 
tallic appearing substance called an alloy. Not all metals 
will mix in this way, and in some cases, definite chemical 
compounds are formed and separate out as the mixture solidi- 
fies, thus destroying the uniform quality of the alloy. In 
general the melting point of the alloy is below the average 
of the melting points of its constituents, and it is sometimes 
lower than any one of them. 

Antimony alloys with many of the metals, and its chief 
commercial use is for such purposes. It imparts to its al- 
loys high density, rather low melting point and the property 
of expanding on solidification. Such an alloy is especially 
useful in type~ founding, where fine lines are to be repro- 
duced on a cast. Type metal consists of antimony, lead and 
tin. Babbitt metal, used for journal bearings in machinery, 
contains the same metals in a different proportion together 
with a small amount of copper. 

Chemical Properties. In chemical properties, anti- 
mony resembles arsenic in many particulars. It forms the 
oxides, SboOg and SboO.^ and in addition SboO^. It com- 
bines with the halogens elements with great energy, burning 
brilliantly in chlorine to form antimony chloride. SbClg. 
Heated on charcoal with the blowpipe, it is oxidized and 
forms a coating of antimony oxide on the charcoal which 
has a characteristic bluish-wdiite color. 



2i6 Inorganic Chemistry. 

Stibine. The gas stibine, SbHg, is formed under 
conditions which are closely parallel to those which produce 
arsine, and it greatly resembles the latter compound, though 
still less stable than it. The mirror of antimony, produced 
as in the Marsh test for arsenic closely resembles the ar- 
senic mirror, but it is more sooty in appearance, and the 
two can be distinguished with great precision by chemical 
tests. 

Acids of Antimony. The oxides, SbgOg, and SbaOB 
are weak acid anhydrides and are capable of forming two 
series of acids corresponding in formula to the acids of 
phosphorus and arsenic. They are much weaker however, 
and are of small practical importance. 

Sulphides of Antimony. Antimony resembles ar- 
senic in that hydrogen sulphide precipitates it as a sulphide 
when conducted into an acidified solution containing an anti- 
mony compound : 

2 SbCl^ + 3 H^S = Sb.Sg + 6 HCl. 
2 SbCl^ + 5 H^S = Sb^Ss + lo HCl. 
The two sulphides of antimony are called the trisulphide and 
the pentasulphide. They are orange colored substances 
when prepared in this way though the mineral stibnite is 
black. 

Metallic Properties of Antimony. The physical 
properties of the elements are those of a metal, and the fact 
that its sulphide is precipitated by hydrogen sulphide shows 
that it acts like a metal in a chemical way. Many other re- 
actions show that antimony is a metal rather than an acid 
forjning element. The compound Sb(OH)3, while able to 
act as a weak acid, is also able to form salts with strong 
acids. When treated with strong hydrochloric acid, anti- 
mony chloride is formed : 

Sb(OH), 4-3 HCl = SbCl,, + 3 H^O. 



The Nitrogen Family. 2iy 

A number of elements act in this same way, their hy- 
droxides acting under some conditions as weak acids, and 
under others as weak bases. 

Hydrolysis of Antimony Chloride. When antimony 
chloride is treated with water it is decomposed into antimony 
hydroxide and hydrochloric acid: 

SbCla+3H.O=Sb(OH)3+3HCl. 

This reaction is a reversible one, and the relative amounts 
of acid and water present determine in which direction and 
to what extent the reaction will go. When excess of water is 
present, it goes as represented in the equation; if excess of 
hydrochloric acid is added to antimony hydroxide, it goes in 
the reverse way. Sometimes the decomposition is not com- 
plete, and a basic salt is formed, thus : 

SbCL+2H20=Sb(OH)2Cl+2HCI. 

The decomposition of a salt by water in either of these 
ways is called hydrolysis. 

BISMUTH. 

Occurrence. Bismuth is usually found in the un- 
combined form in nature, and less frequently as oxide and 
sulphide. Most of the bismuth of commerce comes from 
Saxony, and from Mexico and Colorado, but it is not an 
abundant element, and is rather costly. 

Preparation. It is prepared by merely heating the 
ore containing the native bismuth and allowing the melted 
metal to run out into suitable vessels. 

Physical Properties. Bismuth is a heavy, crystal- 
line, brittle metal nearly the color of silver, but with a slightly 
rosy tint which distinguishes it from other metals. It melts 
at a low temperature (270°), and has a density of 9.8. It 
is not acted upon by the air at ordinary temperatures. 

Chemical Properties. Heated strongly in the air it 
burns to form the oxide 6120^, and before the blowpipe on 
charcoal bismuth eives a coatin-^ of oxide of a vellowish- 



2i8 Inorganic Chemistry. 

brown color which is easily distinguished from the one 
formed by other metals. It combines very readily with the 
halogens elements, powdered bismuth burning readily in 
chlorine. It is not very easily acted on by hydrochloric acid, 
but nitric and sulphuric acids act upon it in the same way 
that they do upon copper. 

Uses. Bismuth finds its chief use as a constituent 
of alloys, particularly in those of low melting point. Some 
of these melt in hot water, as Wood's metal, consisting of 
bismuth, lead, tin and cadmium. 

Compounds of Bismuth. Unlike the other elements 
of this group, bismuth has almost no acid properties. Its 
chief oxide, BigOg, is basic in its properties. It dissolves in 
strong acids and forms salts of bismuth : 

BuO, + 6 HCl = 2 BiCl, + 3 H^O. 
Bi^Og + 6 HNO3 = 2 Bi(NO)3 + 3 H;0. 

When the salts of bismuth are treated with water they 
are partially decomposed, giving basic salts. Thus : 
BiClg + 2 H,0 = Bi(OH)X]+ 2 HCl. 

" Bi(N03)3 + 2H,0 = Bi(6H)oN03 + 2 HNO3. 
When these basic salts are treated with excess of acid, the 
normal salt is once more formed. 

Bi(0H)oN03 + 2 HNO3 = Bi(N03)3 + 2 H,0. 

In this respect the salts of bismuth are similar to anti- 
mony chloride. In the presence of water they undergo hydro- 
lysis, forming basic salts. lii very dilute solution they may 
be completely hydrolyzed into hydroxide and free acid. 

BiCL + H,0 = Bi(OH)3 + 3 HCl. 
The basic nitrate and carbonate of bismuth are used m 
medicine. 



CHAPTER XIX. 

SILICON — BORON. 

Next to oxygen, silicon is the most abundant element. 
It does not occur free in nature, however, but its compounds 
are very abundant and of the greatest importance. It oc- 
curs almost entirely in combination with oxygen as silicon 
dioxide, Si02, often called silica ; or with oxygen and va- 
rious metals in the form of salts of silicic acids, or silicates. 
These compounds form a large fraction of the earth's solid 
crust. Most plants absorb small amounts of silica from the 
soil. It is also found in minute quantities in animal or- 
ganisms. 

Preparation. The element is prepared by reducing 
pure powdered quartz with magnesium powder : 

SiOo + 2 Mg = 2 MgO + Si 
or by a more complicated reaction in which sodium fluo- 
silicate NaoSiFg, is heated with sodium : 

Na.SiF, + 4 Na = 6 NaF + Si. 

Properties. As an element it resembles carbon in 
many respects. It can be obtained in several allotropic forms, 
corresponding to those of carbon. The crystallized form is 
very hard, and inactive toward reagents. The amorphous 
variety has in general properties more similar to charcoal. 

Compounds of Silicon with Hydrogen and the Halo- 
gens. Silicon hydride SiH^, corresponds in formula to 
methane, CH^, but its properties are more like those of 
phosphine, PHg. It is a very inflammable gas of disagree- 
able odor, and as ordinarily prepared, takes fire spontane- 
ously, from the presence of impurities. 

Silicon combines w-ith the elements of the chlorine 
family to form such compounds as SiC^, SiF^. Of these, 

(219) 



220 Inorganic Chemistry. 

silicon fluoride is the most familiar and interesting. As 
stated in the discussion of fluorine it is formed when hydro- 
fluoric acid acts upon any compound of silicon. With silica 
the reaction is thus expressed : 

SiO^ + 4 HF =; SiF, + 2 H^O. 

It is a very volatile, invisible, poisonous gas. In con- 
tact with water it is partially decomposed as shown in the 
equation : 

SiF, + 4 H2O r^ 4 HF + Si(OH),. 

The hydrofluoric acid so formed combines with an ad- 
ditional amount of silicon fluoride forming the complex fluo- 
silicic acid, H2SiFg, 

4 HF + 2 SiF, = 2 H^SiFg. 

Silicides. As the name indicates, silicides are binary 
compounds, consisting of silicon and some other element. 
They are very stable at high temperatures, and are usually 
made by heating the appropriate substances in the electric 
furnace. The most important one is carborundum which is 
a silicide of carbon of the formula CSi. It is made by heat- 
ing coke and sand which is a form of SiOg in an electric 
furnace, the process being extensively carried on at Niagara 
Falls. 

SiOo + 3 C = CSi + 2 CO. 
The substance so prepared, consists of beautiful purplish 
black crystals, which are very hard. Carborundum is used 
as an abrasive, that is to polish and grind very hard sub- 
stances. Ferro-silicon is a silicide of iron, alloyed with an 
excess of iron, which finds extensive use in the manufacture 
of certain kinds of steel. 

Oxide of Silicon. Silicon forms but one oxide, SiOg, 
called the dioxide or silica. This substance is found in a 
great variety of forms in nature, both in the amorphous, and 
in the crystalline condition. In the form of quartz it is found 
in beautifully formed prismatic crystals, sometimes of great 



Silicon-Boron. 221 

size, and when pure, it is perfectly transparent and color- 
less. Some colored varieties are given special names, as 
arriethyst (violet), rose quartz (pale pink), smoky and milky 
quartz (colored and opaque). Other varieties of silicon 
dioxide some of which also contain water are chalcedony, 
onyx, jasper, opal, agate and flint. Sand and sandstone are 
largely silicon dioxide. 

Properties. As obtained by chemical processes, sili- 
con dioxide is an amorphous, white powder. In the crystal- 
lized state, it is very hard, and has a density of 2.6. It is 
insoluble in water, and in most chemical reagents, and re- 
quires the hottest oxyhydrogen flame for fusion. Acids, ex- 
cepting hydrofluoric acid, have little action on it, and it re- 
quires the most energetic reducing agents to deprive it of 
oxygen. It is the anhydride of an acid, and consequently it 
dissolves in fused alkalis to form silicates. Being non- 
volatile, it will drive out most other anhydrides when heated 
to a high temperature with their salts, especially when the 
silicates so form.ed are fusible. The following reactions il- 
lustrate this property : 

Na2C03 + SiO^ = Na.SiOg + CO^. 
• Na^SO^ + SiO^ = NaoSiOg -f SO3. 

These reactions illustrate the law of mass action very 
well. The anhydrides COj and SO3 being gases at the tem- 
perature of the reaction, escape as fast as they are formed. 
Their active mass is therefore kept very small, and the reac- 
tion proceeds to the end instead of reaching an equilibrium. 

Silicic Acids. Silicon forms two simple acids : ortho- 
silicic acid, H^SiO^, and metasilicic acid HgSiO.^. Ortho- 
silicic acid is formed as a jelly-like mass when orthosili- 
cates are digested with strong acids such as HCl. On at- 
tempting to dry this acid, it loses water, passing into meta- 
silicic, or common silicic acid. 

H.SiO, = H.SiO, + H,0. 



222 Inorganic Chemistry. 

Metasilicic acid when heated, breaks up into siUca and 
water thus : 

H^SiOs = H2O + SiO^. 

This reaction recalls the corresponding decomposition 
of carbonic acid into carbon dioxide and water. 
H2CO3 = H^O + CO2. 

Kaolin, Al4(Si04)3, is the aluminium salt of ortho- 
silicic acid, and clay is simply kaolin mixed with other 
mineral substances. 

Polysilicic Acids. Silicon has the power to form 
a great many complex acids which bear somewhat the same 
relation to orthosilicic acid that pyrophosphoric acid bears 
to orthophosphoric acid. Several molecules of the ortho 
acid unite with the loss of water. Thus we have 
3 H.SiO, = H.SigOs + 4 H2O. 

These acids cannot be prepared in pure state, but their 
salts form many of the crystalline rocks in nature. Feld- 
spar for example, has the formula KAlSigOg, and is a mixed 
salt of the acid H^SijOg, whose formation is represented in 
the equation above. Many other examples will be met in the 
study of metals. 

Glass. When sodium and calcium silicates, together 
with silicon dioxide, are heated to a very high temperature, 
the mixture fuses to a transparent liquid which on slowly 
cooling, passes into the solid called glass. Instead of starting 
with sodium and calcium silicates, it is more convenient and 
economical to heat sodium carbonate (or sulphate), and 
lime with an excess of clean sand, the silicates being formed 
during the heating. 

Na^CO;, + SiO^ = Na^SiOg + CO,. 
CaO + SiOg = CaSiOg -f CO^. 

The mixture is heated below the fusing point for some 
time so that the escaping carbon dioxide may not spatter 



. SUicon-Boron. 22:^ 

the hot liquid; the heat is then increased and the mixture 
kept in a state of fusion until all gases formed in the re- 
action have escaped. 

The way in which the melted mixture is handled in the 
glass factory depends upon the character of the article to 
be made. Many articles such as bottles are made by blow- 
ing the plastic glass into hollow moulds of the desired shape. 
Other objects, such as lamp chimneys, are made by getting 
a lump of plastic glass on the end of a hollow iron rod, and 
blowing it into the desired shape without the help of a 
mould, great skill being required in the manipulation of the 
glass. Window glass is made by blowing large hollow cylin- 
ders of glass, about 8 ft. long by 2 ft. in diameter, cutting 
them open longitudinally, softening them in an oven, and 
rolling them into flat plates. Plate glass is cast into flat 
slabs which are then ground and polished to perfectly plane 
surfaces. ^ 

The ingredients mentioned above make a soft, easily 
fusible glass. If potassium carbonate is substituted for the 
sodium carbonate the glass is much harder and less easily 
fused; increasing the amount of sand has somewhat the 
same effect. Potassium glass is much used in making chem- 
ical glassware, since it resists the action of reagents better 
than the softer sodium glass. If lead oxide is substituted 
for the whole or a part of the lime, the glass is very soft, 
but has a high index of refraction and is valuable for 
making optical instruments and artificial jewels. 

Various substances fused along with the glass mix- 
ture give the glass characteristic colors. The amber color 
of common bottles is due to iron compounds in the glass ; 
in other cases iron colors the glass green. Cobalt com- 
pounds color it deep blue, while manganese gives it an 
amethyst tint, and uranium compounds a peculiar yellowish 
green. Since iron is nearly always present in the ingre- 



224 Inorganic Chemistry. 

clients, glass is usually slightly green. This color can be 
removed by adding the proper amount of manganese dioxide, 
for the amethyst color of manganese and the green of. iron 
together produce white light. 

Glass is not a definite chemical compound, and its com- 
position varies between wide limits. It is really a solution of 
various silicates, such as those of calcium and lead, and of 
silicon dioxide in fused sodium or potassium silicate. This 
solution gradually turns into a solid on cooling without the 
dissolved substances separating from the solvent. The com- 
pounds which are used to color the glass are often converted 
into silicates which then dissolve in the glass, giving it a 
uniform color. In other cases, as in the milky glasses which 
resemble porcelain in appearance, the color or opaqueness is 
due to the finely divided color material evenly distributed 
throughout the glass, but not dissolved in it. Milky glass is 
made by melting tin oxide or some similar substance into 
the glass, while copper or gold in metallic form scattered 
through the glass, give it shades of red. 

Carbon and Silicon are thus seen to be the central ele- 
ments in the framework of nature, in the organic and inor- 
ganic worlds respectively. They resemble each other in the 
abundance and complexity of their compounds ; but this 
complexity is due to entirely different causes. The variety 
and complexity of carbon compounds is due to the ability of 
the carbon atoms to unite with each other in chains or rings, 
while with silicon it is due to the formation of complex silicic 
acids. 

BORON. 

Boron is a. somewhat rare element which the periodic 
law assigns to the aluminium family, but which can best be 

studied in connection witli carbon and silicon. 



Silicon-Boron. 225 

Occurrence. Boron is never found free in nature. It 
occurs as boric acid, H3BO3, and in salts of p'olyboric acids, 
which usually have very complicated formulas. 

Preparation and Properties. Boron can be prepared 
from its oxide by reduction with magnesium exactly as in 
the case of silicon. It resembles silicon very strikingly in 
properties, occurring in several allotropic forms, being very 
hard when crystallized, and being rather inactive toward 
reagents. It forms a hydride BHg, and combines directly 
with the elements of the chlorine family. Boron fluoride, 
BF3, is very similar to silicon fluoride in its mode of forma- 
tion and chemical properties. 

Oxides of Boron. Boron forms one well known ox- 
ide, B2O3, called boric anhydride. It is formed as a glassy 
mass by heating boric acid to a high temperature. It ab- 
sorbs water very readily, uniting with it to form boric acid 
again : 

B2O3 + 3 H2O = 2 H3BO3. 
In this respect it differs from silicon dioxide. 

Acids of Boron, Boric acid, H3BO3, is found m 
nature in considerable quantities, and forms one of the chief 
sources of boron compounds. It is found dissolved in the 
water of hot springs in some localities, particularly in Italy. 
Being volatile with steam, the steam which escapes from 
these springs has some boric acid in it. It is easily obtained 
from these sources by condensation and evaporation, the 
necessary heat being supplied by other hot springs. 

Boric acid crystallizes in pearly flakes, which are greasy 
to the touch. In the laboratory the acid is easily prepared 
by treating a strong, hot solution of borax with sulphuric 
acid. Boric acid being sparingly soluble in water,. crystallizes 
out on cooling. 

Na^B.O, + 5 H^O + H^SO, = Na^SO, + 4 H3BO3. 



226 Inorganic Chemistry. 

The substance is a mild antiseptic and is often used on this 
account m medicine, and as a preservative for canned foods 
and in milk. 

Metaboric and Polyboric Acids. When boric acid is 
gently heated, it is converted into metaboric acid, HBOg. 

H3BO3 = HBO2 + H2O. 
On heating metaboric acid to a somewhat higher tem- 
perature, tetraboric acid, H2B4O7 is formed. 

4HBO2 = H2B4O7 + H2O. 
Many other complex acids of boron are known. 

Borax. Borax is the sodium salt of tetraboric acid, 
having the formula Na^B^Oy, 10 HoO. It is found in some 
arid countries, as southern Cahfornia and Thibet, but is now 
made commercially by boiling colmanite (calcium poly- 
borate) with sodium carbonate, calcium carbonate being 
precipitated, and borax crystallizing from the solution. 

When heated, borax at first swells up greatly while the 
water of cystallization is being given off, and then melts to a 
clear glass. This glass has the property of easily dissolving 
many metallic oxides, and on this account borax is used 
as a flux in soldering to clean the metallic surfaces to be 
soldered from the film of oxide with which they are likely 
to be covered. These oxides often give a characteristic color 
to the clear borax glass, and borax beads are therefore often 
used in testing for the presence of metals instead of the 
metaphosphoric acid bead already described. 

The reason that oxides dissolve in borax is that borax 
contains an excess of acid anhydride, as can be more easily 
seen if its formula is written 2 NaBO^ + B2O3. The metal- 
lic oxide combines with this excess of acid anhydride form- 
ing a mixed salt of metaboric acid. 

Borax is used a great deal in domestic ways as a mild 
alkali to replace lye and soap. It is strongly alkaline to 



^lAr^61S4i 



Silicon-Boron. 227 

indicators such as litmus, and its solutions have the proper- 
ties of an alkali. 

' This may seem strange, especially when it is remembered that 
borax contains an excess of acid anhydride. The explanation 
is to be found in the fact that the acids of boron are very 
-weak ones — that is they are very little dissociated into ions. 
In solution the salts of such weak acids are hydrolysed some- 
what, that is, they fall apart into free acid and free base. 

Na2B40T+2H20=2NaOH+H2B407. 
Since the acid is little dissociated, there are few hydrogen 
ions in the solution; the base, on the other hand, freely dis- 
sociates into its ions Na and OH. The hydroxyl ions there- 
fore largely exceed the hydrogen ions in number, and the solu- 
tion has a strong alkaline reaction. 

Borax is also used in the making of enamels, and for 
the manufacture of soaps designed to be used with hard 
water. 



MM 27 1S95 



An Introduction 

TO 

Inorganic Chemistry. 



LIST OF THE ELEMHNTvS. THEIR SYMBOLS 
AND ATOMIC WEIGHTS. 



H = l = 16 

Aluminium ...Al .. 26.9 27.1 

Antimony .....Sb ..119.1 120 

Argon ...A ... 39.6 39.9 

Arsenic As . . 74.4 75 

Barium .Ba ..134.4 137.4 

Bismuth ...... Bi ...206.9 208.5 

Boron ...B .... 10.9 11 

Bromine .Br ... 79.36 79.96 

Cadmium .....Cd ..111.6 112.4 

Caesium ... ....Cs ...132 133 

Calcium Ca .. 39.7 40 

Carbon C ... 11.91 12 

Cerium Ce ..139 140 

Chlorine ......CI ... 35.18 35.45 

Chromium ....Cr .. 51.7 52.1 

Cobalt Co .. 58.56 59 

Columbium ...Cb .. 93.3 94 

Copper ^Cu .. 63.1 63.6 

Erbium ....... E ...164.8 166 

Fluorine ...... F ... 18.9 19 

Gadolinium ...Gd ..155 156 

Gallium ...... Ga .. 69.5 70 

Germanium ...Ge .. 71.5 72 

Glucinum Gl .. 9.03 9.1 

Gold .......... Au ..195.7 197.2. 

Helium He . . 4 4 ' 

Hydrogen ....H ... 1 1.01 

Indium .In ...113.1 114 

Iodine I ....125.9 126.85 

Iridium Ir . ..191.5 193 

Iron Fe . . 55.6 56 

Krypton Kr .. .sl.2 81.8 

Lanthanum ...La ..137 138 

Lead Pb ..205.35 206.9 

Lithium Li ... 6.98 7.03 

Magnesium ...Mg .. 24.18 24.36 

Manganese ... ..Mn .. 54.6 55 

Mercury Hg ..198.8 200.3 



H = l 0==16 

Molybdenum... Mo .. 95.3 96 

Neodymium ...Nd ...142.5 143.6 

Neon Ne .. 19.9 20 

Nickel Ni .. 58.3 58.7 

Nitrogen N ... 13.93 14.04 

Osmium Os ..189.6 191 

Oxygen O ... 15.88 16 

Palladium ....Pd ..105.2 106 

Phosphorus ...P 30.77 31 

Platinum .....Pt ...193.3 194.8 

Potassium . . . .K ... 38.86 39.15 

Praseodymium. Pr ..139.4 140.5 

Rhodium Rh ..102.2 103 

Rubidium .....Rb .. 84.76 85.4 

Ruthenium ....Ru ..100.9 101.7 

S-miarium . . . .Sa .. .148.9 150 

Scandium Sc .. 43.8 44.1 

Selenium Se .. 78.5 79.1 

Silicon Si ... 28.2 28.4 

Silver Ag ..107.12 107.93 

Sodium Na .. 22.88 23.05 

Strontirm . . . . Sr ... 86.94 87.6 

Sn]phur S ... 31.83 32.06 

Tantalum Ta ..181.6 183 

lellnrium Te ..126 127 

Thallium Tl ...202.6 204.1 

Thorium Th ..230.8 232.5 

Thulium Tu ..170 171 

Tin Sn ..117.6 118.5 

Titanium Ti ... 47.7 48.1 

Tungsten .....W ...182.6 184 

Uranium U ...237.7 239.5 

Vanadium ....V ... 50.8 51.2 

Xenon ...X ...127 128 

Ytterbium ....Yt ..172 173 

Yttrium Y ... 88.3 89 

Zinc Zn .. 64.9 65.4 

Zirconium ....Zr ... 90 90.7 



